A Unifying Framework for Understanding Electrooxidation of Small Organic Molecules

A UNIFYING FRAMEWORK FOR UNDERSTANDING
THE ELECTROOXIDATION OF SMALL ORGANIC
MOLECULES FOR FUEL CELL APPLICATIONS
CHENG CHIN HSIEN
(B. Eng. (Hons.), NUS)
A THESIS SUBMITTED
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
DEPARTMENT OF CHEMICAL & BIOMOLECULAR
ENGINEERING
NATIONAL UNIVERSITY OF SINGAPORE
(2011/2012)

I
ACKNOWLEDGEMENT
First and foremost, I would like to acknowledge my thesis supervisor, Professor Lee Jim
Yang, for his support and guidance throughout the course of this project. His sharing on
technical knowledge, advice on my writing skill, and patience in revisions of my thesis,
are the keys for me to deliver this thesis work.
I would like to thank my colleagues in research group, Dr. Liu Bo, Dr. Yang Jin Hua, Dr.
Zhang Qing Bo, Dr. David Julius, Mr. Chia Zhi Wen, Miss Yu Yue, Miss Lu Mei Hua,
for the discussion and help throughout my work and their valuable comments to this
thesis as fellow scientists.
I am thankful for the research scholarship from National University of Singapore, and the
assistance from the technical and administrational staffs of Department of Chemical and
Biomolecular Engineering.
Last but not the least; I would like to thank my family for their forever understanding and
support.

II
TABLES OF CONTENTS
ACKNOWLEGEMENT I
TABLE OF CONTENTS II
SUMMARY XI
LIST OF SCHEMES XIV
LIST OF TABLES XVI
LIST OF FIGURES XVIII
LIST OF SYMBOLS XXIII
CHAPTER 1 INTRODUCTION 1
1.1 Background and Objective 1
1.2 Fuel Cell Fundamentals 2
1.2.1 Basic Fuel Cell Construction 2
1.2.2 Fuel Cell Reactions at Equilibrium 4
1.2.2.1 Thermodynamic Cell Potential at Standard Conditions 4
1.2.2.2 Standard Hydrogen Electrode (SHE) 4
1.2.2.3 Nernst Equation and Reversible Hydrogen Electrode (RHE) 6

III
1.2.3 Fuel Cell Reactions at Non-Equilibrium 7
1.2.3.1 Overpotential and Internal Resistance 7
1.2.3.2 Voltammetry and Current Density 9
1.3 Reconciliation Process 10
1.4 The Capability of Proposed Unifying Mechanism and its Core Principles 12
1.4.1 Different Systems Examined in this Thesis 12
1.4.2 Core Principles for Deducing Unifying Mechanism Framework 13
1.5 Thesis Structure and Comparisons between Current and Proposed
Mechanisms
15
Chapter 1 – Supporting Information 1S 21
1S1 Experimental 21
CHAPTER 2 MAJOR REACTION PATHWAYS IN THE
ELECTROOXIDATION OF SMALL OXYGENATES ON
PLATINUM IN ACIDS
24
2.1 Introduction 24
2.2 The Proposed Unifying Mechanistic Framework 26
2.2.1 Unifying Attributes: Pt&α-C, Pt&O, and Pt&H Interactions 26
2.2.2 CO Adsorption and Electrooxidation 28
2.2.3 HCOOH Adsorption and Electrooxidation 29
2.2.3.1 Dependence of Reaction Pathways on Pt&α-C, Pt&O, and Pt&H
Interactions
31
2.2.3.2 Observations of Surface Geometry Dependency 33
2.2.4 Aldehyde Adsorption and Electrooxidation 34

IV
2.2.4.1 Major Difference between H2C(OH)2/H2CO and HCOOH
Electrooxidations
35
2.2.4.2 Similarities between H2C(OH)2 and HCOOH Electrooxidations 37
2.2.4.3 Comparison between CH3CHO and H2CO 38
2.2.5 Alcohol Adsorption and Electrooxidation 39
2.2.5.1 The Pathways Determined by Pt&α-C and Pt&O Interactions 39
2.2.5.2 Optimization of Surface Geometry and Operating Temperature 41
2.3 Conclusion 43
2S1 Pt&O and Pt&H (*H, *H2O, H2O*) Interactions at 0.4V 45
2S2 Pt&α-C, Pt&O Interactions at 0.4V and around *OH Onset Potentials 47
2S3 Suppression of *CO Formation and Optimization of the Direct
*COOH Pathway when Adsorption as *COOH is Least Affected by
H* and *O-species
48
2S4 Observations of *OCHO* as an Inhibiting Species at High Potentials 50
2S5 *CHO as one of Surface Blocking Species 53
2S6 Conversion of :CROH to *CRO 54
2S7 Stronger Surface Inhibition by CH3CHO than by H2CO 54
2S8 Direct O-Addition Pathways in the Oxidation of Alcohols to
Carboxylic Acids and Hydrated Aldehydes
55
2S9 Selectivity for *CO and *CRO Formation during Alcohol
Electrooxidation and Its Dependence on Step Density
57
2S10 Optimal (110) Step Density for Current Generation 58

V
2S11 Elevated Temperature Enhanced Dehydration 59
2S12 Doubts in Recent Publications Supporting *OCHO* as Reactive
Intermediate
59
CHAPTER 3 COMPLETE ELECTROOXIDATION OF ETHANOL AND
ACETALDEHYDE IN ACIDS AT HIGH POTENTIALS
VIA ADSORBED CARBOXYLATES ON PLATINUM
62
3.1 Introduction 62
3.2 Proposed Mechanisms for the Complete Oxidation of Ethanol and
Acetaldehyde
65
3.3 Supporting Evidence for the Proposed Origin of the second CO2 Peak 67
3.4 Conclusion 68
3S1 Protracted *CO Electrooxidation in the Presence of Adsorbed Acetate 70
3S2 Evidence for *OC(CH3)O* Electrooxidation 72
3S3 The Central Region of the second CO2 Peak via *OCHO* and
*O*OCCO*O*
75
CHAPTER 4 THE INHIBITION OF PLATINUM SURFACE BY
ACETALDEHYDE AND ACETIC ACID FORMATION
DURING ETHANOL ELECTROOXIDATION IN ACIDS
83
4.1 Introduction 83
4.2 Results and Discussion 86
4.2.1 Electrooxidation of CH3COOH and CH3CHO 86

VI
4.2.2 Electrooxidation of Ethanol with CH3CHO or CH3COOH 87
4.2.3 Effects of Pt/C Loading Per Electrode Surface Area 90
4.2.3.1 Overall Activity 91
4.2.3.2 CO2 Efficiency 93
4.2.4 The Appropriateness of the If / Ib Ratio as an Indicator of Catalyst
Tolerance
97
4.3 Conclusion 98
4S1 CH3CHO electrooxidation at various concentrations 100
4S2 Effects of CH3CHO and CH3COOH Addition on Ethanol
Electrooxidation in Different Potential Regions
102
4S3 Observations that Support Direct O-Addition of Alcohol as the Major
Current Contributor in the Reverse Scan
104
CHAPTER 5 PROMOTION OF THE DIRECT O-ADDITION
PATHWAYS IN ALCOHOL ELECTROOXIDATION ON
BIMETALLIC PLATINUM-RUTHENIUM CATALYSTS
106
5.1 Introduction 106
5.2 Results and Discussion 108
5.2.1 Observations Supporting the Enhancement of the Direct O-addition
Pathways
108
5.2.2 Observations of Activation and Deactivation of PtRu Catalysts 111
5.2.3 Proposed Mechanism of PtRu Activation and Deactivation 112
5.2.4 Adverse Effect of Excessive *OH and O* on PtRu Activity 113

VII
5.2.5 PtRu Activation by Cyclic Voltammetric Pretreatment in C2H5OH
between 0.06V and 1.17V
115
5.3 Conclusion 118
5S1 The Activation of Deactivation of PtRu during Methanol
Electrooxidation
119
CHAPTER 6 EFFECTS OF TIN IN PLATINUM-TIN CATALAYSTS
FOR ELECTROOXIDATION IN ACIDS
120
6.2 Review, Results and Discussion 121
6.2.1 Distribution of Sn and Its Effects on CO, Formaldehyde and Methanol
Electrooxidation
121
6.2.2 Similarity between Methanol and Ethanol 123
6.2.2.1 Enhancement of the O-Addition Pathway for Alcohols by *OH on
Sn/SnOx and Weaker Pt&α-C Interaction
123
6.2.2.2 Adsorption is Rate-limiting on Pt-Sn Alloys 124
6.2.3 Difference Between Ethanol and Methanol 125
6.2.3.1 Easier Adsorption of Ethanol 125
6.2.3.2 Inhibition by *OC(CH3)O* during Ethanol Electrooxidation 125
6.2.4 Comparison between Pt-Sn and Pt-Ru 127
6.3 Conclusion 128
6S1 Temperature Effect on the Optimal Sn Distribution for Ethanol 130

VIII
Electrooxidation
CHAPTER 7 HIGH SELECTIVITY OF PALLADIUM CATALYSTS
FOR THE DIRECT DEHYDROGENATION PATHWAY IN
FORMIC ACID ELECTROOXIDATION IN ACIDS
132
7.2 Discussion 133
7.2.1 Strong Pd&H Interaction Results in Weak Pd&O Interaction 134
7.2.2 The Interaction between H* and *CO 135
7.2.3 Enhanced Selectivity for the Direct HCOOH Pathway 138
7.2.4 Optimization of Pd-Based Catalysts for HCOOH Electrooxidation 141
7.3 Conclusion 143
CHAPTER 8 EFFECTS OF IONIZATION ON ETHANOL
ELECTROOXIDATION ON PLATINUM AND
PALLADIUM IN ALKALINE SOLUTIONS
144
8.2 Core Concepts 146
8.2.1 C2 Pathways 146
8.2.1.1 C2 Pathways on Pt 147
8.2.1.2 C2 Pathways on Pd 151
8.2.1.3 Effects of Catalyst Loading per Electrode Surface Area on C2
Pathways
153
8.2.2 Electrooxidation of CO 154
8.2.3 C1 Pathways on Pt and Other Poisoning Species 154

IX
8.2.4 Principles for Catalyst Optimization 156
8.2.4.1 Optimization of Pt Catalysts 156
8.2.4.2 Optimization of Pd-Based Catalysts 156
8.3 Conclusion 157
CHAPTER 8 – Supporting Information 8S 159
8S1 Weaker Effects of Acetic Acid on Pt catalysis under Strongly Alkaline
Conditions
159
8S1.1 Effects of Acetic Acid in the Absence of Ethanol 159
8S1.2 Effects of Acetic Acid in the Presence of Ethanol 162
8S2 Weaker Effects of Acetaldehyde on Pt Catalysis in Strongly Alkaline
Solutions
163
8S3 Formation of CH3CHO by the Direct Dehydrogenation of CH3CH2O-
165
8S4 Effects of Acetaldehyde and pH on Pd Catalysis in Strongly Alkaline
Solutions
166
8S4.1 Acetaldehyde at pH 13.93 167
8S4.2 Acetaldehyde at pH 13.40 169
8S5 Effect of Catalyst Loading per Electrode Area 171
8S6 Pt Surface Dependent Activities and Deactivation Rates 174
CHAPTER 9 CONCLUSIONS AND RECOMMENDATIONS 177
9.1 An Unifying Mechanistic Framework of Reactions 178
9.1.1 Pt-Based Catalysis 178
9.1.2 Pd-Based Catalysis 181
9.2 Considerations for Reactions in a Strongly Alkaline Environment 181

X
9.3 Useful Practical Information for Catalyst Development 182
9.3.1 Comparison of Catalyst Activities 182
9.3.2 Important Indicators from Cyclic Voltammetry 182
9.4 Recommendations 183
REFERENCE 186

XI
SUMMARY
This thesis aims to develop a comprehensive understanding of the electrooxidation of
small oxygenates1
for fuel cell applications, which can satisfactorily explain many of the
experimental observations spanning over a diverse range of catalysts and operating
conditions. This is by reconciling the many disagreements in the current literature on
reaction mechanisms, and infilling the knowledge gaps between systems with different
combinations of catalysts, fuels and operating conditions, together with our own
experimental supporting evidences.
With such a unifying understanding for various systems, one can predict the catalyst
performance and provide the guidelines for a practical catalyst design for the specific fuel
molecule. A cross comparison between various fuels with understanding on the predicted
limit of improved catalyst design, could further help in selecting the best choice of fuel
from the anode reaction perspective. This is important since the current bottleneck in
portable fuel cell development is on the anode electrooxidation reaction.
The systems which were analyzed in this thesis are representative of low temperature fuel
cell operations and include the following variables
1
Oxygenates in this thesis are with broad definition, i.e. oxygen containing compounds from incomplete
oxidation of hydrocarbon molecules

XII
Operating conditions: potential, acidic and strong alkaline solution, temperature,
catalyst loading per electrode surface area
Catalysts: Monometallic Pt with different surface geometries, bimetallic Pt-Ru and Pt-Sn,
monometallic Pd.
Small Oxygenate Molecules: CO, HCOOH, H2CO and its hydrated form H2C(OH)2,
CH3CHO and CH3CH(OH)2, CH3COOH, CH3OH, CH3CH2OH.
For monometallic Pt in acidic condition, for example, current density per unit Pt mass can
be improved by suppressing the formation of surface blocking *CO or *CRO. This can
be achieved by inhibiting C-OH bond cleavage on α-C, or by promoting the addition of
C-OH bond to α-C. This in turns requires the weakening of Pt&α-C interaction and the
availability of *OH at low potentials. For bimetallic catalysts (e.g. Pt-Ru or Pt-Sn) which
are designed to provide such functionalities, the Pt&α-C interaction has to be optimized
to prevent the over-weakening of the Pt&α-C interaction which can turn the
dehydrogenative adsorption of oxygenate into a rate limiting step (e.g. in alcohol
electrooxidation). The electrooxidation of C2 molecules is more complex since C-C bond
cleavage and adsorbed acetate (*OC(CH3)O*) inhibition are additional considerations.
Strongly alkaline condition is able to weaken both the *C(CH3)O and *OC(CH3)O*
inhibition, and improves the catalyst activity. Strongly alkaline condition could even help
the C-C bond cleavage on Pt, it is however not a perfect solution since large inhibiting
molecules via aldol reaction could gradually deactivates the catalyst. The optimization of
the catalyst design and operating conditions can in principle be based on the tuning of

XIII
two fundamental attributes: 1) the interaction between the catalytic site and adsorbed *H,
*C-species and *O-species; 2) the equilibrium between (and among) adsorbed species
and dissolved species (e.g. RCHO  RCH(OH)2, RCOOH *OCRO*). However, these
two attributes may be mutually compensating in the electrooxidation of more complex
molecules. Therefore, from a practical perspective, HCOOH may be the best fuel for
portable applications.

XIV
LIST OF SCHEMES
Scheme 2.1 The proposed general reaction scheme for HCOOH
electrooxidation. The direct dehydrogenation pathway (CO2
formation via *COOH) is the most desirable for current generation.
It occurs when the surface is not blocked by *CO and is most
favorable when adsorption as *COOH is least interfered by H* and
*O-species (i.e. at around ptzc). T*CO formation can be minimized
by a weaker Pt&α-C interaction; and by the competing adsorption
of species in the blue boxes. Once T*CO is formed, it can only be
removed effectively by oxidation when T*OH becomes abundant
(i.e. at high V, via the pathway in red).
30
Scheme 2.2 A proposed general reaction scheme for H2C(OH)2 electrooxidation.
It is analogous to HCOOH oxidation in the following aspects: direct
dehydrogenation pathways via O-H cleavage(s) in solution to form
HCOOH and CO2, indirect pathways via surface catalyzed C-OH
cleavage forming inhibiting *CHO and subsequently *CO. The
main difference is the added possibility of *CHO formation from
H2CO, which makes surface inhibition an easier process.
36
Scheme 2.3 Proposed reaction scheme for alcohol electrooxidation illustrating
the direct O-addition pathways to form carboxylic acid or hydrated
aldehyde, and the formation of inhibiting *CRO and *CO species.
The presence of adjacent S*OH at low potentials and an optimized
Pt-C bond strength for desorption are required for high activity
towards direct O-addition pathways.
40
Scheme 3.1 The proposed pathways (non-elementary steps) for the complete
oxidation of C2H5OH and CH3CHO to CO2 in different potential
regions.
66
Scheme 4.1 Suggested reaction scheme for C2H5OH electrooxidation. R is CH3.
Adsorbed species in blue compete for adsorption through Pt-C
mainly on the *T sites. Adsorbed species in red compete for
adsorption through Pt-O mainly at high potentials or on *S sites at
low potentials. Pathways with green, purple, or red arrows require
reaction with *OH and are therefore inhibited by the red adsorbed
species. The difficulty of *OH addition increases from green to
purple to red colored pathways. The *T sites, on the other hand, are
easily passivated by *CRO and *CO at low potentials. Increase in
catalyst loading enhances the re-adsorption of RCHO to *CRO and
suppresses the direct O-addition pathway to RCOOH and
RCH(OH)2 formation (thick green arrow), resulting in higher CO2
selectivity. However, increase in potential normally decreases CO2
selectivity by the preferentially catalyzing the oxidation of
96

XV
RCH2OH and RCHO to RCOOH than the C-C cleavage of *CRO.
However, when a very high catalyst loading is used, re-adsorption
of RCOOH as *OCRO* occurs to suppress the O-addition pathways
colored in green and in purple to rates close to the red colored
pathways. Increase in potential and *OH coverage will therefore
ease the electrooxidation of the red colored *O-carbon residue
species to CO2, improving activity and CO2 selectivity
simultaneously.
Scheme 5.1 Possible changes in the catalyst surface structure during PtRu
(Pt:Ru = 1:1) activation by the cyclic voltammetric treatment with
1.17V anodic scan limit in C2H5OH. Red spheres: Ru. Small brown
spheres: O or OH. Blue spheres: Pt. In PtRu alloys, Pt could be
heavily affected by more adjacent Ru atoms to slow the alcohol
adsorption. Grey spheres: Pt with Pt&C and Pt&O interactions
similar to those in monometallic Pt, to restitute good alcohol
adsorption while keeping the supply of adjacent *OH groups. (How
the specific cyclic voltammetric treatment could modify the PtRu
surface will be explained in §5.2.5).
113
Scheme 8.1 The reaction mechanism from reference for ethanol
electrooxidation. Solid arrows are the reaction pathways at low pH,
while dashed arrows are the pathways for high pH. The
deprotonation of the CH3 group of CH3CHO forms the enolate
anion, CH2=CHO-
with good delocalization of the acquired negative
charge.
145
Scheme 8.2 Proposed reaction mechanism for ethanol electrooxidation on Pt.
The pathways in the lower section enclosed by the red box are
electrooxidation in acidic solutions which has been discussed in
Chapters 2-4. Ionization in strongly alkaline solutions opens up the
pathways in the upper section. Green arrows: reactions with S*OH
at practical anode potentials. Orange arrows: reactions with *OH at
high potentials. Purple arrows: formation of *CRO or :CRO-
on
sites with strong Pt&C interaction (e.g. (110)*T). The adjacent sites
should have moderately strong Pt&O interactions if C-OH cleavage
is involved.
149

XVI
LIST OF TABLES
Table 1.1 Calculation of the ΔG0
and E0
for reactions 4 and 5. (ΔGf
0
: Standard
Gibbs free energy of formation of compounds.
5
Table 1.2 A simple example of deriving a unifying mechanism through the
reconciliation of observations from different but related systems.
10
Table 1.3 Comparison between Current and Proposed Mechanisms 18
Table 2.1 Effects of Pt surface geometry on Pt&α-C, Pt&O, Pt&H interactions
at ~ 0.4V.
27
Table 2.2 The important potentials in 0.1M HClO4, and species from H2O
dissociation that compete with *C-species for adsorption.
28
Table 2S.1 The dominant adsorbed species on Pt basal planes in 0.1M HClO4 46
Table 3.1 Summary of the ethanol reaction mechanisms showing the effects of
Pt&O, Pt&α-C, Pt&β-C interactions on various electrooxidation
pathways.
68
Table 3S.1 Deduction of possible β-C1 adsorbed species for the values of n =
3.7, α-C/β-C = 0.5 measured at Ead = 0.6V
77
Table 4.1 Comparison of inhibiting species in CH3OH and C2H5OH
electrooxidation
93
Table 4S.1 Observations and explanations of voltammetric response in
CH3CHO electrooxidation
102
Table 4S.2 Effects of CH3CHO and CH3COOH Addition on Ethanol
Electrooxidation
103
Table 5.1 Effects of Cyclic Voltammetric Pretreatments on C2H5OH
Electrooxidation and Reduction of Surface Species in 0.1M HClO4
116
Table 6.1 Distribution of Sn/SnOx among the Pt atoms on the catalyst surface
and their effects on the electrooxidation of CO, aldehydes, and
alcohols*
122
Table 6S.1 Product distribution in the effluent of single cell tests at 90˚C 131
Table 8.1 Comparison of processes that affect the rate of ethanol
electrooxidation on Pt under acidic and strongly alkaline conditions
148

XVII
Table 8.2 Comparison of ethanol electrooxidation under acidic and strongly
alkaline conditions on Pd
151
Table 8S.1 Effects of CH3CHO addition to ethanol electrooxidation at pH 13.93
and 13.40
171
Table 9.1 Summary of Pt-catalyzed electrooxidation of different oxygenates
in acidic solutions at room temperature
179
Table 9.2 Summary of the effects of different Pt-based catalysts and operating
conditions
180

XVIII
LIST OF FIGURES
Fig. 1.1 The basic components of a PEMFC. 3
Fig. 2.1 The surface geometry of Pt(100), Pt(111), Pt(110), and a plane with
(110) steps on (111) terraces (i.e. Pt(S)[(n-1)(111)x(110)],
representing (n-1) rows of atoms on (111) terraces before a (110) step.
In this Fig, n = 3). Pt(110) is the plane with maximum (110) step
density on (111) terraces. *T on grey-colored atoms includes the
Pt(100)*T, the Pt(111)*T and the Pt(111)-like *T sites on Pt(S)[(n-
1)(111)x(110)]. *T on orange-colored atoms includes the Pt(110)*T
and the Pt(110)-like *T sites on Pt(S)[(n-1)(111)x(110)].
27
Fig. 2.2 A concave surface with (111) terraces and (110) step hollow sites *S
(red triangles) but without the (110)-like *T sites.
42
Fig. 2S.1 Plot of CO-coverage on Pt(111) and Pt(100) surfaces in CO-free 0.1
M H2SO4 as a function of the dosing potential (squares). The total
charge without double layer correction (triangles), calculated from the
hydrogen adsorption region of the voltammogram, is also included.
48
Fig. 2S.2 Cyclic voltammograms for two Pt basal planes in 0.1 M HCOOH +
0.1 M HClO4. The solid lines represent first potential scans starting at
50 mV vs RHE. Dotted lines correspond to the voltammogram in an
electrolyte without HCOOH. Insets: enlarged voltammograms in
selected potential regions; units, mAcm-2
. Scan rate 50 mV/s.
49
Fig. 2S.3 Cyclic voltammogram for a 12
CO-covered Pt electrode in 0.5 M
H2SO4+ 0.1 M H13
COOH at a sweep rate of 50 mV/s; and the
corresponding plot of the integrated band intensities of *12
CO and
*O13
CHO*in the positive-going scan (solid line). The dotted line
represents the oxidative removal of a *12
CO monolayer in an
electrolyte without H13
COOH.
51
Fig. 2S.4 Potential oscillations observed in 0.5 M H2SO4+ 0.1 M formaldehyde
at the applied current of 10 mA on a Pt film electrode and the
corresponding plot of integrated band intensities of T*CO, :CO, and
adsorbed formate in the 18s-35s time frame.
53
Fig. 2S.5 CVs of Pt single crystals in 0.5 M CH3OH and 0.5 M HClO4 at a scan
rate of 2mV/s: (a) Pt basal planes, (b) Pt surfaces with (110) steps on
(111) terraces.
58

XIX
Fig. 2S.6 a) Current transient for a double-potential step from 0.05 to 0.9 (2 s)
and then to 0.6 V (vs. RHE) in 10 mM HCOOH with 0.5M H2SO4. b,
c) Transients of the integrated band intensities of COL, COB, formate,
and (bi)sulfate taken from a set of time-resolved IR spectra of the Pt
electrode surface collected simultaneously with the current transient at
80 ms intervals.
60
Fig. 2S.7 Curve fitting for derived correlation “iformate α k (θformate)3/2
/ cHCOOH ”. 61
Fig. 3.1 DEMS mass intensities of 12
CO2 and 13
CO2 during oxidative stripping
of adsorbed residues from isotopically labelled ethanol (a),
acetaldehyde (b) on Pt. Stripping was carried out with and without
pre-reductive stripping in the hydrogen adsorption region (a), or at
different adsorption potentials (Ead) (b).
64
Fig 3S.1. (a) Integrated IR intensities of *CO from pre-adsorbed CO and
C2H5OH residues; (b) SEIRAS spectra of the oxidation of C2H5OH
residues at different potentials. Electrolyte: 0.1M HClO4. Ead = −0.1V
Ag/AgCl ~ 0.16V RHE.
71
Fig. 3S.2 Cyclic voltammograms of an E-TEK catalyst (20µg Pt /cm2
) in 0.1M
HClO4 with different CH3COOH concentrations at 100mV/s after
stabilizing pre-scans in HClO4 (a: to 1.17V, b: to 1.47V). See text for
the description of regions (1) to (5).
73
Fig. 3S.3 DEMS mass intensities of CO2 formation from the oxidative stripping
of 1-propanol (a), iso-propanol (b), and four butanol isomers (c-f) pre-
adsorbed at various potentials.
78
Fig. 3S.4 Cyclic voltammograms of an E-TEK catalyst 20µg Pt /cm2
in 0.1M
HClO4 with different oxalic acid concentrations at 100mV/s after
stabilizing pre-scans in HClO4.
79
Fig. 4.1 Steady state cyclic voltammograms of E-TEK catalyst @ 5µg Pt /cm2
in 1M C2H5OH at 10mV/s: (a) stationery electrode vs rotating disc
electrode @ 1000rpm; (b-c) rotating disc electrode @ 1000rpm in the
presence of different CH3CHO or CH3COOH concentrations. The
insets in (b-c) show the percentage current remaining after the addition
of CH3CHO or CH3COOH (the arrows indicate scan directions). The
different potential regions of interest as demarcated by vertical black
lines are discussed in Supporting Information 4S2.
88
Fig. 4.2 Steady state cyclic voltammograms of the electrooxidation of 1M
C2H5OH (or 1M CH3OH) in 0.1M HClO4 on an E-TEK Pt/C catalyst.
The catalyst loading on a stationary electrode was varied to give
different Pt weights per electrode area. The inset shows the percentage
91

XX
current density in C2H5OH electrooxidation relative to the base case
loading of 5µg Pt/cm2
.
Fig. 4.3 Total oxidation charge and the percentage of which from CO2
production at different potentials during the chronoamperometry of
C2H5OH electrooxidation on a 4mgPt/cm2
loaded carbon paper in a
stationary electrolyte system at room temperature. (experimental
details in Chapter 1 )
95
Fig. 4S.1 Cyclic voltammograms of a E-TEK Pt/C catalyst with 20µg Pt /cm2
loading in HClO4 solutions with different acetaldehyde (AAld)
concentrations at 100mV/s. (a) 1st
cycle after holding at 0.05V for 30s;
(b) stabilized response; (c) stabilized response on a 1000rpm rotating
disc electrode. The voltammograms have been corrected for the
background current in 0.1M HClO4. Prior to this the catalyst was
scanned repeatedly in HClO4 until a stable response was established.
101
Fig. 5.1 (A) Steady-state cyclic voltammograms of electrooxidation of ethanol
and acetaldehyde on Pt (E-Tek 20wt%) and PtRu (E-Tek 20wt%). (B)
Cyclic voltammograms of electrooxidation of 1M ethanol from 1st
to
35th
scans (inset: forward scan current in the 0.4V-0.5V region). All
measurements were taken in 0.1M HClO4 at 10mV/s
110
Fig. 5.2 Cyclic voltammogram of 1M ethanol electrooxidation in 0.1M HClO4 on 5µg
PtRu /cm2
with an anodic scan limit of 0.7V. Scan rate: 10mV/s
112
Fig. 5.3 Cyclic voltammograms in 0.1M HClO4 on 20µg PtRu and 20µg Pt
/cm2
. For PtRu (A) shows the 1st
scans with different anodic potential
limits (0.7V and 1.17V) without any pretreatment; and (B) shows the
1st
scans with anodic potential limit of 1.17V after different
pretreatments: (blue - 35 scans to 1.17V in 0.1M HClO4 only, cyan &
pink - 35 scans to 0.7V (cyan) & 1.17V (pink) in 1M C2H5OH + 0.1M
HClO4). For Pt steady state response is used for both (A) and (B). Scan
rate: 100mV/s
115
Fig. 5S.1 Cyclic voltammogram of 1M methanol electrooxidation in 0.1M HClO4 on
10µg PtRu /cm2
from 1st
to 60th
scans. Scan rate: 10 mV/s. Activation:
increase in current density at potentials below 0.6V from scan 1 to scan 35,
the increase in peak current density is more persistent, until scan 60.
Deactivation: decrease in current density at potentials below 0.6V from scan
35 onwards.
119
Fig. 6.1 The 20th
scan cyclic voltammograms of the electrooxidation of 1M
C2H5OH in 0.1M HClO4 on an E-TEK Pt3Sn/C catalyst at 10mV/s in
the presence of different concentrations of extraneously introduced
CH3CHO (A) or CH3COOH (B). The catalyst loading was 5µg
Pt3Sn/cm2
. Argon was continuously purged to eliminate dissolved O2
127

XXI
and to generate turbulence to improve external diffusion of products.
Fig. 7.1 Cyclic voltammograms of Pd electrode in 0.5 M H2SO4 at 20mV/s. (–
) after 600s of CO adsorption at 0.40V and then 600s at 0.00V without
CO in the solution; (--) voltammogram obtained after complete
oxidation of CO adsorption products.
136
Fig. 7.2 The effect of adsorption potential of CO on charge passed to the
electrode during CO adsorption (Qads) and during subsequent
electrooxidation of *CO (QOx).
137
Fig. 7.3 Voltammograms of the electrooxidation of formic acid on Pd(111) and
Pd(100) in 0.1 M HClO4 containing 0.1 M formic acid. Scanning rate:
20 mV/s
140
Fig. 7.4 Voltammograms of formic acid electrooxidation on modified Pd
catalysts in 0.5 M H2SO4 containing 0.5 M formic acid at 50 mV/s:
(A) comparison between Pd/C, Pd/RT (rutile TiO2) and Pd/CMRT
(carbon modified rutile TiO2); (B) comparison between Pd/C, Pt/C,
alloyed Pd20Pt, and Pt decorated Pd/C (Pd:Pt = 20:1).
142
Fig. 8.1 Stabilized cyclic voltammograms of ethanol electrooxidation in 0.1M
HClO4 (with/without rotation at 1000rpm) and in 0.85M KOH.
Catalyst loading: 5µg Pt/cm2
. Scan rate: 10mV/s.
147
Fig. 8.2 Effect of acetic acid (A) and acetaldehyde (B) addition on stabilized
cyclic voltammograms (CVs) of ethanol electrooxidation in 0.1M
HClO4 (with 1000rpm rotation). Catalyst load: 5µg Pt/cm2
. Scan rate:
10mV/s. This Figure shows the decrease in j/V slope due to a slower
direct O-addition reaction caused by species competing with S*OH
(A). The right shift in the j-V curves is caused by *C(CH3)O which
interferes with ethanol adsorption (B). These voltammetric responses
should be compared with the responses sown in Fig. 8.1 and Fig. 8.3.
150
Fig. 8.3 Steady state voltammograms of ethanol electrooxidation in 0.85M
KOH electrolyte with different CH3COO-
concentrations. Catalyst
loading: 5µg Pd/cm2
. Scan rate: 10mV/s.
153
Fig. 8S.1 Effect of pH on the voltammogram of Pt(111) at 30 mV/s in (a) 20
mM CH3COOH + 0.02, 0.1, or 0.3 M HClO4 at pH (I) 0.7 (—); (II)
1.1 (---) and (III) 1.9 (···); and (b) mixtures of CH3COOH and
CH3COOK (total concentration = 0.2M) with pH (I) 5.1 (---), (II) 5.6
(···), and (III) 6.0 (—), respectively (from, with the potential scale
converted to SHE (bottom) and RHE (top)).
160
Fig. 8S.2 Voltammograms of ethanol oxidation on E-TECK Pt/C (20µg Pt /
cm2
) in 0.85M KOH with the addition of different amounts of
161

XXII
CH3COOH which leads to 0.85M K+
and final CH3COO-
concentrations as indicated. Scan rate = 100mV/s. Inset:
voltammograms in 0.1M HClO4, for comparison
Fig. 8S.3 Effects of CH3COOH addition on the stabilized voltammograms of
C2H5OH electrooxidation on E-TEK Pt/C (5µg/cm2
) in alkaline
solutions, at 10mV/s. For clarity of presentation, only the forward
scans are shown for CH3COOH addition.
163
Fig. 8S.4 Effect of CH3CHO addition to stabilized voltammograms of C2H5OH
electrooxidation on E-TEK Pt/C (5µg/cm2
) in alkaline solutions, at
10mV/s. The inset shows the blocking effect of *C(CH3)O from
CH3CHO in acidic solutions for comparison.
164
Fig. 8S.5 Linear sweep voltammograms of (a) alcohols (10mM) with high j, (b)
alcohols (10mM) with low j on Au electrode in 0.1 M NaOH (pH =
13) with a scan rate of 50 mV/s (a-b, value in bracket is pKa); (c) plots
of the onset potential versus the pKa (value in bracket is pKa and
onset potential); and (d) Tafel plots of the corresponding alcohols.
166
Fig. 8S.6 Effect of acetaldehyde addition on the voltammogram of E-TEK Pd/C
(5µg/cm2
) in 0.85M KOH at 10mV/s, in the absence (A) and presence
of in 1M ethanol (B). The small spike around 0.16V in (B) occurred at
the instant CH3CHO was added at the end of the reverse scan of the
“before adding CH3CHO” voltammogram.
168
Fig. 8S.7 Effects of acetaldehyde addition on the voltammogram of E-TEK
Pd/C (5µg/cm2
) in a solution containing 0.25M OH-
(0.85M K+
and
0.60M CH3COO-
) at pH 13.40 at 10mV/s, in the presence of in 1M
ethanol.
170
Fig. 8S.8 Effect of E-TEK Pd/C loading per electrode surface on the
voltammogram of 0.85M KOH + 1M ethanol at 10mV/s. Current
density is normalized by a) Pd mass, or b) electrode geometrical area.
172
Fig. 8S.9 Effect of E-TEK Pt/C loading per electrode surface on the
voltammogram of 0.85M KOH + 1M ethanol at 10mV/s in. Current
density is normalized by a) Pd mass, or b) electrode geometrical area.
173
Fig.8S.10 Voltammograms of ethanol electrooxidation on Pt(111) and Pt(110)
for the 1st
(a) and the 20th
(b) cycle in 0.5M ethanol and 0.1M NaOH,
at 10mV/s.
175

XXIII
LIST OF SYMBLES
Symbols regarding surface sites and species involved in reactions
* a general adsorption site when there is no need to be specific about the site
geometry
*S step hollow site
*T terrace top site
: bridge binding site (terrace bridge site)
triple binding site (terrace hollow site)
*H adsorbed H
*C-species adsorbed species with C atom bound to the surface
*O-species adsorbed species with O atom bound to the surface
*O-carbon
residue
adsorbed carbon residue with O atom bound to the surface (e.g. *OCH3). It
is more specific than *O-species since it excludes *OH and O*.
R a H atom or an alkyl group, if it appears in a chemical formula, e.g.
RCOOH representing carboxylic acid
–H* surface catalyzed dehydrogenation
H+
proton (hydronium ion, H3O+
, is sometimes written as H+
for
simplification).
e-
electron
–H+
–e-
a proton release from adsorbed species via interactions with surrounding
H2O or OH-
with the simultaneous transfer of an electron to the electrode
Symbols regarding calculations involved potentials and current (density)
E thermodynamic potential at equilibrium (unit in volt, V)
E0
thermodynamic potential at equilibrium at standard conditions (V)
E0
cell thermodynamic cell potential at standard conditions (V)

XXIV
η overpotential (V)
V applied or operating potential (V)
V cell operating cell potential (V)
I current (A) = dQ/dt
Q charge (C)
t time (s)
r internal resistance ( )
T Temperature (K, or ˚C)
F Faraday constant 96485 C/mol e-
R gas constant 8.314 (J K−1
mol−1
)
ΔG0
Gibbs free energy changes per mole of reaction (J/mol reaction) at standard
conditions
aox chemical activity of oxidized form of a redox species
ared chemical activity of reduced form of a redox species
z the number of electrons exchanged per mole of reaction (mol e-
/ mol
reaction)
n number of moles of reactant
dn/dt the moles of reactant converted per time
mcat mass of metal catalyst (mg)
ECSA electrochemical surface area (cm2
)
GEA geometrical electrode area (cm2
)
Loadcat catalyst metal loading per geometrical electrode area (mg metal catalyst /
cm2
electrode)
J current density (A/cm2
ECSA)

XXV
SHE Standard Hydrogen Electrode
RHE Reversible Hydrogen Electrode
CE Coulombic Efficiency
EE Energy Efficiency

Chapter 1
1
CHAPTER 1
INTRODUCTION
1.1 Background and Objective
Fuel cells are able to convert the stored chemical energy in fuel molecules directly into
electricity by spatially separating the electrooxidation of fuel and the electroreduction of
oxygen. As heat is not involved as an intermediate step, electricity generation by fuel
cells is not subjected to the Carnot limit as in the case of heat engines. Hence fuel cells
can be used at relatively low temperatures (e.g. ambient temperature) providing on-
demand electricity so long as there is fuel in the system and the fuel cell circuit is closed.
Fuel cells therefore have an inherent advantage over rechargeable batteries which require
mains power and substantial recharge time to replenish the depleted charge. However,
fuel cells also have their fair share of technical challenges such as storage and delivery of
fuel especially if the latter is a gas (e.g. hydrogen) and the use of (expensive) catalysts.
While the use of liquid fuels can alleviate the fuel storage problem, liquid fuels are also
more difficult to electrooxidize than hydrogen, resulting in low power density and low
energy conversion efficiency.

Chapter 1
2
The bottleneck in direct liquid fuel cells 2
is the poor performance of fuel electrooxidation
at low temperatures. Technological breakthrough is possible only if better catalysts are
available for our choices of fuel molecules and operating conditions (e.g. temperature and
pH). The traditional empirical approach of exploring statistically many different catalysts
and evaluating their performance under different combinations of fuel molecules and
operating conditions is hardly efficient. An in-depth understanding of the reaction
mechanisms, on the other hand, will be more useful to guide the catalyst design and to
anticipate the limitations in different fuel molecules and different operating conditions.
However, most of the work done up to date has targeted at specific catalyst-fuel-
operating condition combinations and as such is of limited utility to derive any general
understanding if the results are examined in isolation without reference to other related
studies. Hence there is no lack of “conflicting theories” in the literature. The objective of
this thesis is therefore to seek a unifying understanding of the reaction mechanisms for
the electrooxidation of small oxygenates (mainly C1-C2 alcohols, aldehydes and
carboxylic acids) to explain satisfactorily most of the experimental observations in the
literature and all of the original results in the thesis study.
1.2 Fuel Cell Fundamentals
1.2.1 Basic Fuel Cell Construction
2
Direct Liquid Fuel Cells: Fuel cells that convert the chemical energy in liquid fuel directly into electricity,
without an intermediate steam reforming process to convert the liquid fuel to hydrogen.

Chapter 1
3
The basic elements of a fuel cell and fuel cell principles are summarily described in this
section before the discussion of reaction mechanisms. A typical fuel cell consists of an
anode, a cathode, an external circuit to conduct the electrons, and an electrolyte in the
interior of the fuel cell between the electrodes to conduct either H+
or OH-
.
For example, in a hydrogen proton exchange membrane fuel cell (PEMFC) (Fig. 1.1), H2
is electrooxidized at the anode. The e-
and H+
formed in the oxidation reaction are
transported from the anode to the cathode through the external circuit and the proton
exchange membrane respectively. The e-
arriving at the cathode then combines with the
oxygen there to form H2O.
Fig. 1.1. The basic components of a PEMFC.

Chapter 1
4
1.2.2 Fuel Cell Reactions at Equilibrium
1.2.2.1 Thermodynamic Cell Potential at Standard Conditions
Thermodynamics determines the energy released in a redox reaction. For a 100%
conversion of this energy into electricity in fuel cells, the half-cell reactions on both
electrodes have to be at equilibrium. The difference between the equilibrium electrode
potentials of the cathode and the anode is therefore the maximum cell potential possible.
The following is an example illustrated with H2 as the fuel.
O2 + 4H+
+ 4e-
2H2O (1.229V S.H.E.) (1)
2H+
+ 2e-
 H2 (0V S.H.E.) (2)
2H2 + O2 2H2O (E0
cell=1.229 – 0 = 1.229V) (3)
where S.H.E is the acronym for the standard hydrogen electrode (vide infra), and E0
cell is
the thermodynamic cell potential at standard conditions. Reaction 1 is the cathode
reaction (O2 electroreduction), reaction 2 is written as the reverse of the anode reaction
(H2 electrooxidation, by convention electrode reactions are often written as reduction
reactions), and reaction 3 is the overall fuel cell reaction.
1.2.2.2 Standard Hydrogen Electrode (SHE)
SHE is often used as the reference for which other equilibrium electrode potentials are
quoted. 0V SHE refers to the equilibrium potential of 1 bar H2 in a 1M [H+
] (pH = 0)
electrolyte over a platinum black surface at 25˚C. The equilibrium electrode potentials of

Chapter 1
5
other fuel molecules at standard conditions can be calculated from the Gibbs free energy
changes of half-cell reactions by the following thermodynamic relationship:
E0
= -ΔG0
/ zF
where E0
and ΔG0
are the equilibrium electrode potential (V) and Gibbs free energy
changes per mole of reaction (J/mol reaction) at standard conditions respectively; z is the
number of electrons exchanged per mole of reaction (mol e-
/ mol reaction), and F is the
Faraday constant 96485 C/mol e-
.
An example calculation of the ΔG0
and E0
for the reduction of CO2 to ethanol (reaction 4)
is shown in Table 1.1. Such calculations are important to determine the equilibrium
electrode potentials of different fuel molecules. The potential of a full cell reaction
(reaction 5) can also be determined similarly.
2CO2 + 12H+
+ 12e-
C2H5OH + 3H2O (4)
C2H5OH + 3O2  2CO2 + 3H2O (5)
Table 1.1 Calculation of the ΔG0
and E0
for reactions 4 and 5. (ΔGf
0
: Standard Gibbs free
energy of formation of compounds, from [1])
Compound CO2 (g) H2O (l) C2H5OH (l)
ΔGf
0
(kJ/mol) -394.4 -237.1 -174.8
ΔG0
4 (kJ/mol) (-174.8) + 3(-237.1) – 2(-394.4) = -97.3
E0
4 -(-97.3)(1000)/12/96485 = 0.084V
ΔG0
5 (kJ/mol) 2(-394.4) + 3(-237.1) – (-174.8) = -1325.3
E0
5 -(-1325.3)(1000)/12/96485 = 1.145V
Alternatively, E0
1 (=1.229) - E0
4 = 1.145V

Chapter 1
6
1.2.2.3 Nernst Equation and Reversible Hydrogen Electrode (RHE)
SHE is defined with respect to a fixed set of conditions (1M, pH0, 1bar and 25˚C). The
equilibrium electrode potentials at other conditions can be calculated from the Nernst
equation.
Electrode reaction: ox + e ↔ red
E= E0
+ (RT/zF) ln(aox/ ared)
where R is the gas constant 8.314 (J K−1
mol−1
), T is the temperature (K), aox and ared are
the chemical activity of oxidized and reduced forms of the redox species. For the
reduction of H+
to hydrogen in aqueous solution, aox and ared can be approximated by the
pressure of gaseous hydrogen in bar and the H+
concentration in M respectively.
Since H+
is always involved in the electroreduction reactions investigated in this study,
the prevailing equilibrium potential is a function of the solution pH. With 1 unit increase
in the pH (~1 order of magnitude lower in [H+
]), the equilibrium potential would decrease
by ~ (8.314 x 298 / 96485) ln(1/10) = 0.0591V (59.1 mV).
Nevertheless, since the equilibrium potentials of oxygen and fuel molecules all involve
the participation of H+
, changes in pH occur to the same extent on both electrodes and
hence do not affect the overall cell potential. The reversible hydrogen electrode (RHE) is

Chapter 1
7
another reference equilibrium electrode. It is defined with respect to the electrolyte in use
rather than a 1M [H+
] (pH = 0) standard solution. It is more convenient for the
comparison of electrochemical reaction rates at different pH and is the de facto reference
electrode to use in this study unless stated otherwise. The relation between SHE and
RHE is the following:
RHE = SHE – 0.0591 (pH)
1.2.3 Fuel Cell Reactions at Non-Equilibrium
For practical fuel cell operations, neither the fuel electrooxidation reaction at the anode
nor the oxygen electroreduction at the cathode is at equilibrium. A finite reaction rate is
the result of a sufficient number of reactant molecules overcoming the barrier to reactions
at conditions away from the equilibrium in each half cell. Hence the reaction rate would
depend on the reactant and product concentrations, and the impetus provided to surmount
the barriers to reactions. For electrochemical reactions, this impetus can be delivered as
heat or applied potential. Therefore, reaction rate depends on temperature and on how far
the applied potential is away from the equilibrium electrode potential.
1.2.3.1 Overpotential and Internal Resistance
Overpotential (η) is defined as the difference between the applied potential (V) and the
equilibrium potential (E) of a half-cell reaction.
η = V – E

Chapter 1
8
Overpotential is present at both anode and cathode as the impetus to overcome the
barriers against the activation of redox species and the diffusion of reactant and product
species between the electrode surface and the solution bulk. The “activation overpotential”
is high in the presence of strongly adsorbed species on the catalyst surface because
additional driving force is needed to remove these species by reaction and/or by
desorption. Besides, the transport of ions (e.g. H+
) through the electrolyte also has to
overcome the barrier due to the solution internal resistance (r). The operating fuel cell
voltage (Vcell) is therefore the thermodynamic cell potential reduced by the sum of the
overpotentials and the product of internal resistance and current (I).
Vcell = Ecell - | η | anode - | η | cathode – I.r
For instance in a direct ethanol fuel cell (DEFC), if ethanol electrooxidation at the anode
occurs at 0.7V and oxygen electroreduction at the cathode occurs at 0.8V; the anode and
cathode overpotentials are 0.7 – 0.084 ~ 0.616V, and 0.8-1.229 = -0.429V respectively
(Table 2.1). The overall overpotential of this fuel cell is therefore |0.616| + |-0.429| =
1.045V. The operating full cell voltage will hence be 1.145 – 1.045 (or 0.8 – 0.7) – I.r =
0.1V – I.r In this example, overpotentials deplete about 90% of the equilibrium cell
potential, leaving only ~10% for use under practical conditions. An effective catalyst is
one which could reduce the overpotential as much as possible.

Chapter 1
9
1.2.3.2 Voltammetry and Current Density
Voltammetry, or measurements of the current response to a linearly varying potential, is a
standard electroanalytical technique for assessing the reactivity of an electrochemical
half-cell reaction. A significant current flow at low overpotentials is an indication of
satisfactory activation by an effective catalyst on the electrode.
The measured current is usually normalized by the electrochemical surface area (ECSA)
of the catalyst to yield a measure of the intrinsic activity of the surface sites, or by the
mass of the precious metal in the catalyst to indicate metal utilization, or simply by the
electrode surface area to give a nominal current density if there is no need to emphasize
either of the above.
These current densities are intensive quantities that are measures of reaction rates with
different emphasis. The inter-conversion between them is shown below.
J = I / ECSA = (dQ/dt) / ECSA = zF(dn/dt) / ECSA (i.e. the form of reaction rate)
I / ECSA = I / [mcat . (ECSA / mcat)] = I / [GEA . Loadcat . (ECSA / mcat)]
where J is the current density (A/cm2
ECSA); I is the current (A) or the charge transfer
per time, dQ/dt; Q is the charge (C); t is time (s); n is the number of moles of reactant
and dn/dt is the moles of reactant converted per time; mcat is the mass of metal catalyst
(mg); GEA is the geometrical electrode area (cm2
); Loadcat is the catalyst metal loading
per geometrical electrode area (mg metal catalyst / cm2
electrode).

Chapter 1
10
1.3 Reconciliation Process
Table 1.2 A simple example of deriving an unifying mechanism through the
reconciliation of observations from different but related systems.
Observations Description Mechanisms based on
Simple Deductions
Exceptions
1 Fish lives in water Whatever lives in water is
fish
Many
exceptions
2 Fish has no limbs
with digits
Whatever has no limbs with
digits is fish
Many
exceptions
3 Fish has spine Whatever has spine is fish Many
exceptions
4 Fish lives in water,
has spine, and has no
limbs with digits
Aquatic vertebrates that lack
limbs with digits is fish
Dolphin,
tortoise, whale,
hagfish, etc
5 Fish breathes by gill Whatever breathes by gill is
fish
Tadpole
Reconciliation: Mechanisms 1-3 with their many exceptions clearly indicate their
inadequacy as a unifying mechanism. The exercise also highlights the inadequacy of
using the information in observations 1-3 in isolation for formulating the unifying
mechanism. Mechanism 4 is the reconciliation of mechanisms of 1-3. There are many
observations (numerous species of fishes) which support mechanisms 4 & 5, making
either of them appear to be correct. However, some exceptions are revealed after careful
examination. The unifying mechanism below hence comes from careful reconciliation of
mechanisms 4-5
Reconciled Mechanism: Fish is aquatic vertebrate (or craniate) animal that respires by
gill and lacks limbs with digits, even when it is matured.
Table 1.2. illustrates a reconciliation process using fish as a simple example. Any single
observation from 1 to 5 in Table 1.2, is insufficient to deduce a unifying mechanism. A
reconciliation process that considers as much as various observations is hence needed.

Chapter 1
11
Similarly, a mechanism proposed in a research paper may be based on insufficient
experimental observations or limited scopes of study. The uniqueness of the mechanism
is also not assured since there could be other mechanisms which are consistent with the
same set of (limited) observations. A unifying mechanism, on the other hand, has the
ability to explain as many observations as possible in different but related systems. In
general, the more observations that could be explained by the unifying mechanism, the
stronger is the consistency and confidence level of the mechanistic understanding. On the
other hand, the unifying mechanism is a reconstruction exercise based on the clues drawn
from disparate sources of related observations and information similar to solving the
mystery of a detective case.
The reconciliation process or reconstruction exercise requires sophisticated analysis to as
many observations as possible. For a detective, a lot of effort has to be spent in finding
clues and analyzing them before claiming those clues as evidences. Similarly, to
construct a unifying understanding on the electrooxidation of small oxygenates, it is very
important to carefully analyze whatever observations reported in literature, since it is
possible to deduce a different explanation based a same set of experimental observations,
and we need to analyze which explanation can be better linked to other observations. A
detective may sometimes design a “trap” to let the criminal to reveal himself, so as in this
thesis we do have our own experiments (Supporting Information S1) to prove certain
concepts. However, we would like to highlight that published experimental observations

Chapter 1
12
in literature are taken as important as our own experimental observations, since all the
observations have to be analyzed in order to develop a unifying understanding framework.
In some cases, a careful analysis over published experimental observations may even
eliminate the need to conduct our own experiment.
1.4 The Capability of Proposed Unifying Mechanism and its Core Principles
1.4.1 Different Systems Examined in this Thesis
The variables in fuel cell reactions can first be organized into different categories by the
type of fuel molecules used, the catalyst(s) involved and the operating conditions. Each
category is then expanded into subcategories for different specific situations. There is
therefore an almost infinite number of possible combinations that can be examined. This
thesis will only look at the most representative systems over a sufficient variety of fuel-
catalyst-operating condition combinations, as shown below.
Fuel molecule: CO, HCOOH, H2CO and its hydrate H2C(OH)2, CH3CHO and its hydrate
CH3CH(OH)2, CH3COOH, HOOCCOOH, CH3OH and CH3CH2OH
Catalyst: monometallic Pt with different surface geometries, catalyst loading per unit
electrode surface area (Chapters 2, 3, 4, 8), bimetallic Pt (Pt-Ru in Chapter 5 and Pt-Sn in
Chapter 6) and monometallic Pd (Chapter 7, 8).
Operating condition: potential, pH (acidic in Chapters 2-7, alkaline in Chapter 8),
temperature

Chapter 1
13
Currently, there is no general mechanistic framework that can rationalize or reconcile the
multitude of observations under such a wide variety of reaction systems. This is the
unique contribution of this thesis project.
1.4.2 Core Principles for Deducing Unifying Mechanism Framework
To construct a building, beams and pillars are needed to strengthen the structure.
Similarly, to construct a unifying mechanism framework, core principles are needed to
link up observations over various fuel-catalyst-operating condition combinations. There
are two core principles being applied over this thesis:
I. interactions between the catalytic site and adsorbed *H, *C-species and *O-species
II. interactions between and among adsorbed species and dissolved species
(In this thesis, * is used to represent a general adsorption site when there is no need to be
specific about the site geometry. The adsorbate atom which is bound to surface site is
identified next to the * symbol.)
The first core principle can be used to explain the effect of various catalyst geometries
and the distribution of second metal (oxide) to the adsorption rate and selectivity among
various adsorbed species, which will subsequently affect the overall reaction rate and
reaction selectivity. The interactions between the catalytic site and adsorbed *H, *C-
species and *O-species are also influenced by electrode potential. For example, a higher
potential facilitates interaction to *O-species, e.g. the formation of *OH and *O.

Chapter 1
14
The second core principle can be further categorized into i) interactions between
adsorbed species, e.g. oxidation of adsorbed intermediate by *OH; ii) interactions
between dissolved species, e.g. the equilibrium concentration ratio between hydrated and
unhydrated aldehyde (RCH(OH)2  RCHO, R could be a H or an alkyl group); iii)
interaction between adsorbed and dissolved species, e.g. a strongly adsorbed *CO and
*CRO will block the adsorption of other species from the solution.
To further zoom into these three categories, we would like to highlight some simple but
important concepts which are first time suggested (or at least uncommon in literature):
1) Between strongly adsorbed and weakly (or unstably) adsorbed intermediates
requiring oxidation by *OH, the weakly (or unstably) adsorbed intermediate is
easier to be oxidized. This leads to impact to reaction selectivity in alcohol
electrooxidation.
2) Comparing CH3CH(OH)2  CH3CHO to H2C(OH)2  H2CO, the acetaldehyde
has a much higher equilibrium [RCHO] / [RCH(OH)2] concentration ratio, thus
acetaldehyde is much easier to be adsorbed into *CRO as compared to
formaldehyde. Similarly, a higher temperature enhance the dehydration (e.g.
RCH(OH)2 to RCHO), and hence the *CRO formation is also facilitated.
3) The role of *OH is not only in oxidation of other reaction intermediate, it also
affects the reactant adsorption and the formation of certain critical transition
intermediate. This is the major cause of hysteresis between forward and backward

Chapter 1
15
scan during cyclic voltammetry (as well as other tests with step-up or step-down
potential change). For example for HCOOH electrooxidation on Pt(100), the
surface is inactivated by blocking *CO during forward scan, while with remained
*OH during the backward scan, the reaction favors direct pathway. Furthermore,
the addition of Ruthenium (Ru), tin (Sn), and their oxides provides *OH for
alcohol oxidation at lower potential during forward scan but retain the *OH on
adjacent Pt affecting alcohol adsorption during backward scan.
4) With increase pH, the dissolved and adsorbed species will shift their equilibrium
towards anion forms, e.g. RCOOH  RCOO-
. This will shift the adsorption into
*OCRO* to a higher R.H.E. potential. A very high pH could even turn surface
blocking *CRO into :CRO-
, which we believe to be an reactive intermediate (“:”
represents bridge binding to two atoms). These are another examples how a
dissolved species (e.g. increase pH by higher cation concentration) influences
other dissolved and adsorbed species, and hence the reactions being affected.
From the above brief discussion, various fuel-catalyst-operating condition combinations
are well linked by the two core principles.
1.5 Thesis Structure and Comparisons between Current and Proposed Mechanisms
This thesis is set out in 9 Chapters. In the following Chapters, Chapter 2 is focused on
development of a unifying framework for understanding the electrooxidation of formic
acid, C1-C2 aldehydes and alcohols on Pt in acidic condition (only pathways not more

Chapter 1
16
difficult than *CO oxidation are covered); and Chapter 3 investigates reaction pathways
for the complete oxidation of ethanol and acetaldehyde to CO2 at high potentials, and
proposes concern for catalyst design for ethanol electrooxidation. In Chapter 4, inhibition
effect of acetaldehyde and acetic acid during ethanol electrooxidation on Pt is examined.
In Chapter 5 and 6, the effect of added Ru/RuOx and Sn/SnOx to Pt and their distribution
are investigated, and the various effects of *OH in alcohol electrooxidation is thoroughly
discussed. Chapter 7 is focused on HCOOH electrooxidation on Palladium (Pd) in acidic
condition, with a same unifying understanding as in Chapter 2 is applied. Chapter 8 is
devoted to an examination of effects of ionization on ethanol electrooxidation on Pt and
Pd in strong alkaline solutions, as an extension from the unifying mechanism from acidic
condition. Finally, Chapter 9 concludes this research work with summary table
highlighting the major findings of this thesis work.
Due to the many conflicting mechanisms in the literature, this thesis will not have a
specific chapter on literature review. Instead our proposed unifying mechanisms will be
introduced first, followed by the presentation of experimental evidence (drawing from the
literature and some of our own) supporting the various ramifications of the unifying
mechanisms in different chapters. Often the most important observations and major
arguments are given in the chapter main pages, with secondary observations and
additional arguments in the Supporting Information of each Chapter. This thesis structure
aims to facilitate the readers’ appreciation of how the unifying mechanism framework
may be applied to different conditions and situations, without the burden of information
overload. This particular way of organizing the information may however make it less

Chapter 1
17
easy to identify the difference between current hypotheses in the research area and the
mechanisms proposed in this thesis. In order to reduce such potential compromise and to
more clearly differentiate the contributions of this thesis from previous research, Table
1.3 is provided as a checklist and roadmap for comparing between current and proposed
mechanisms. It is recommended that the readers refer to this Table after completing each
chapter as a summary of the major findings therein.
Among the findings in this checklist, readers may like to pay special attention to our
proposed direct O-addition pathways for alcohol electrooxidation (in red fonts in Table
1.3) since they are the major pathways for current generation in the unifying mechanism
for alcohols. Besides, their repeated occurrence in a wide range of conditions (including
Pt with or without Ru or Sn addition in acidic conditions; and Pt and Pd in strongly
alkaline conditions) also adds credence to the acceptance of these proposed pathways.
In Table 1.3, “*S”, “*T” are two specific types of adsorption sites, namely the step hollow
site and the terrace top site respectively. The geometry of these sites is illustrated in
Fig.2.1 (Chapter 2.) For simplicity the balance of * is omitted in some equations. “–H*”
represents surface catalyzed dehydrogenation, “–H+
–e-
” represents a proton release via
interactions with surrounding H2O or OH-
with the simultaneous transfer of an electron
to the electrode.

Chapter 1
18
Table 1.3. Comparison between Current and Proposed Mechanisms
Chapter 2 (Reaction Pathways on Pt at Practical Anode Potentials)
Fuel
Molecule
Current Mechanisms Proposed Mechanisms Desorbed
ProductReactive
Intermediate
Blocking
Intermediate
Reactive
Intermediate
Blocking
Intermediate
HCOOH
*COOH *CO T*COOH
(*OCHO* is not
reactive but
suppresses *CO
formation)
T*CO (*OCHO*
only weakly
suppresses the
*COOH
pathway)
CO2
*OCHO* *CO
*COOH, *CO, *OCHO*
H2CO /
H2C(OH)2
*OCH2O* *CO T*CH(OH)2
:C(OH)2
T*CO,
T*CHO
HCOOH
CO2
CH3CHO /
CH3CH(OH)2
Non *CO,
*C(CH3)O
Non T*CO,
T*C(CH3)O
Non
CH3OH
:CHOH + *OH
 *CH(OH)2
*CO
:CHOH + S*OH
 T*CH(OH)2
T*CO,
T*CHO
HCOOH
*CH2OH
(unknown
pathway)
T*CH2OH +
S*OH 
H2C(OH)2
H2C(OH)2
CH3O* –
H*H2CO
CH3CH2OH
*C(CH3)O +
*OH 
CH3COOH
*CO
:C(CH3)OH +
S*OH 
T*C(CH3)(OH)2
T*CO,
T*C(CH3)O
CH3COOH
:C(CH3)OH +
*OH 
*C(CH3)(OH)2
CH3CH2O* –
H*  CH3CHO T*CH(CH3)OH
+ S*OH 
CH3CH(OH)2
CH3CH(OH)2
*CH(CH3)OH –
H*  CH3CHO
Chapter 3 (Reaction Pathways at Potentials Higher than *CO Electrooxidation on Pt)
Fuel
Molecule
Current Mechanisms Proposed Mechanisms Desorbed
ProductIntermediate at
High Potentials
Conflicting
Findings
Intermediate at
High Potentials
Remark
CH3CH2OH,
CH3CHO *CHx
*CHx was found
to convert to
*CO at low
potentials
*O-species, e.g.
*OCHO*,
*O*OCCO*O*,
*C(CH3)O*
They suppress
*OH formation,
but their
oxidation
require *OH
CO2

Chapter 1
19
Chapter 4 (The Effects of *OC(CH3)O* and *C(CH3)O Blocking and Catalyst loading)
Current Proposed
No detailed study on the catalyst poisoning effect of
CH3COOH and CH3CHO during CH3CH2OH
electrooxidation
*OC(CH3)O* competes with *OH formation;
T*C(CH3)O blocks reaction sites and suppresses
CH3CH2OH adsorption.
No detailed study on the effect of catalyst loading
per unit electrode surface area
Catalyst loading affects CH3COOH and CH3CHO
diffusion, and therefore the Pt activity
If / Ib ratio was thought to indicate the catalyst
tolerance to poisoning
If / Ib ratio is not a proper indicator to measure the
catalyst tolerance to poisoning.
Chapter 5 - 6 (Effects of Ru & Sn Addition to Pt)
Current Proposed
Enhanced *CO oxidation by
i) (common reason): bi-functional effect with Ru or
Sn sites providing *OH
ii) (common reason): (Surface) electronic ligand
effect by weakening Pt-CO but strengthening Pt-OH
iii) (unpopular reason): Sn-OH weakens adjacent Pt-
CO by intermolecular interaction
All three factors are valid, but their most important
effect is not enhancing *CO oxidation, but
suppressing *CO formation and facilitating alcohol
electrooxidation via
:CROH + *OH  T*CR(OH)2 RCOOH + H+
+ e-
T*CHROH + *OH  RCH(OH)2
The intermolecular interaction with *OH is very
important and worthy of more attention.
Chapter 7 (HCOOH Electrooxidation)
Current Proposed
*COOH  CO2, but lacks detailed explanation on
good selectivity over *CO formation
With the proposed *CO formation mechanism in
Chapter 2, and the well-known strong Pd&H
interaction, the good selectivity is explained.
Chapter 8 (Electrooxidation in Strongly Alkaline Solutions)
Fuel
Molecule
Current Mechanisms Proposed Mechanisms (for Pt)
Reactive Pathway Reactive Pathway Desorbed
Product
CH3CH2OH /
CH3CH2O-
*C(CH3)O + *OH  CH3COOH
:C(CH3)O-
+ S*OH 
T*C(CH3)(OH)O-
(negligible surface inhibition by
*C(CH3)O)
CH3COO-
For Pd, with to stronger Pd&H but
relatively weaker Pd&C:
:C(CH3)O-
+ *OH-

T*C(CH3)(OH)O-
+ e-
CH3COO-

Chapter 1
20
CH3CH2O* – H*  CH3CHO T*CH(CH3)O-
+ S*OH 
CH3CH(OH)O-
CH3CH(OH)O-
*CH(CH3)OH – H*  CH3CHO T*CH(CH3)O-
 CH3CHO + e-
CH3CHO
CH3CH(OH)O-
/ CH3CH(OH)2
/ CH3CHO /
CH2CHO-
*C(CH3)O + *OH  CH3COOH CH3CH(OH)O-
– H* 
T*C(CH3)(OH)O-
(very active pathway on Pd)
CH3COO-
CH2CHO-
 H2C**CHO-
 ? CH2CHO-
 H2C**CHO-
? CO3
-
It is reasonable that the reader may feel unconvinced by this unifying mechanism, due to
the unavoidable conflicts to other proposed mechanisms in literature. However, readers
are strongly encouraged to investigate and analyze the reported experimental
observations (in literature and in this thesis work) and propose your own unifying
mechanism to cover the same range of various reaction systems as in this thesis. A better
buy-in of the unifying mechanistic framework in this thesis may only come after the
reader experiences through a similar thinking process. Nevertheless, we are open to the
possibility of a different unifying understanding.

Chapter 1
21
CHAPTER 1 – SUPPORTING INFORMATION 1S
1S1 Experimental
The experiments in this thesis study were mainly voltammetric measurements of
commercial catalysts under different conditions. They were carried out following the
experimental details shown below.
1) Materials
Analytical grade (99.8% minimum) methanol, ethanol, and acetic acid from Merck and
99.5% acetaldehyde from Fluka were used. Commercial E-Tek 20wt% Pt/C, PtRu/C, or
Pt3Sn/C catalysts were used for the electrooxidation of methanol or ethanol in acidic or
alkaline electrolyte solutions. The acidic electrolyte is prepared by diluting concentrated
HClO4 (70% in water, Sigma-Aldrich) with Millipore pure water; and the alkaline
electrolyte is prepared by dissolving analysis grade KOH pellet (Merck) into Millipore
pure water.
2) Preparation of Catalyst Ink and Electrode

Chapter 1
22
10mg catalyst powder was first dispersed in 5ml of ethanol by ultrasonication for at least
30 minutes, followed by the addition of 5ml pure water and 50µl of 5wt% Nafion
solution (Aldrich) under ultra-sonication. For voltammetric measurements, the catalyst
ink was dispensed onto a 5mm diameter (~0.2 cm2
) glassy carbon working electrode to a
loading that ranged from 5 to 320 µg Pt /cm2
. For CO2 selectivity measurements, a
sufficient change in CO2 concentration in effluent gas was needed, a working electrode
with a high catalyst loading (4mgPt/cm2
) on carbon paper (1.5cm2
) was therefore used.
3) Voltammetry
A Pt foil and a Ag/AgCl (3M) reference electrode were used as the counter electrode and
reference electrode in a standard three-electrode cell setup. The reference electrode was
connected through a Luggin capillary to minimize the I.r drop. All recorded potentials
were converted to the RHE scale. The electrolyte was 0.1M HClO4 (for an acidic reaction
environment) or 0.85M KOH (for a strongly alkaline reaction environment) with
controlled concentrations of CH3COOH or CH3CHO in the absence or presence of
C2H5OH. Cyclic voltammetry (CV) was carried out at room temperature. The
voltammograms shown are stabilized responses unless indicated otherwise.
4) CO2 Selectivity Measurements (for Chapter 4 only)
For the measurements of CO2 selectivity, the C2H5OH solution was purged with a
constant flow of Argon (30 ml/min) and the downstream gas was monitored by a Telaire

Chapter 1
23
7001 CO2 Monitor. The amount of CO2 corresponding to 5 min of chronoamperometric
C2H5OH electrooxidation on working electrode was recorded. Prior to use the CO2
detector was calibrated by CO2 generated from the injection of a predetermined amount
of NaHCO3 into an excess of HClO4 solution.

Chapter 2
24
CHAPTER 2
MAJOR REACTION PATHWAYS IN THE
ELECTROOXIDATION OF SMALL OXYGENATES ON
PLATINUM IN ACIDS
2.1. Introduction
Platinum is the most common catalytic metal component for the electroreduction of
oxygen and the electrooxidation of alcohols, which are respectively the cathode and
anode reactions of a direct alcohol fuel cell. A good mechanistic understanding of these
reactions on Pt is a scientific undertaking of practical importance as it can lead to more
effective catalyst designs. The direct alcohol fuel cells, despite their potential for
converting the chemical energy in renewable fuels such as wood methanol and bioethanol
directly into electricity, are beset with significant challenges most notably the poor
performance of the catalyst for the complete electrooxidation of alcohol molecules to
CO2.
Despite extensive basic research on the catalysis of specific alcohols and their partial
oxidation products, e.g. aldehydes, carboxylic acids and CO, there has been little effort in
analyzing the common features in the electrooxidation of these different but related
compounds, and clarifying the effects of catalyst surface structure and applied potential.
For example, in formic acid electrooxidation, adsorbed formate (*OCHO*) has been

Chapter 2
25
attributed as a catalyst poison [2, 3], an active intermediate in the direct pathway [4-10],
and even a catalyst [11]. There are also many papers which do not consider adsorbed
formate at all. With the debates on the dual pathway (i.e. electrooxidation through
*COOH and *CO (traditional) or *OCHO* and *CO [4, 5, 7-10]) and triple pathway
(electrooxidation through *COOH, *CO and *OCHO* [2]) of HCOOH electrooxidation
remain unsettled3
, the mechanisms for more complex electrooxidation reactions involving
methanol, formaldehyde, ethanol, or acetaldehyde, are often developed without the latest
understanding of HCOOH electrooxidation. After a careful and systematic survey of the
literature, we found that studies on a particular compound can actually bring new insights
to understanding the reactions of other related compounds.
This chapter puts forward an unifying understanding of the adsorption and
electrooxidation of CO, HCOOH, H2C(OH)2 (hydrated H2CO), CH3CH(OH)2 (hydrated
CH3CHO), CH3OH and C2H5OH. The unifying understanding was derived after
painstakingly analyzing a large amount of published data for underlying patterns and
trends. Several new perspectives and reactions are proposed for the first time to fill the
knowledge gap and to reconcile some inconsistent explanations in the literature. A good
reaction mechanism should have general utility, and is able to explain most (if not all) of
the reported observations Hence the unifying mechanism is based on the consolidation
3
Although most recent publications tend to suggest *OCHO* to be reactive intermediate, their
observations and analysis will be discussed in Supporting Information 2S12 to show that it is not yet
conclusive.

Chapter 2
26
and reconciliation of many independent observations and is not derived from a few
isolated or discrete observations.
For pedagogical reasons, the unifying framework will be presented first, without the finer
details, to explain the effects of reactants, adsorbed species, surface geometry and applied
potential. This is purposely done to demonstrate the generality and the deductive power
of the unifying framework. The observations which were used to construct the unifying
framework, i.e. the supporting evidence, are mostly relegated to Supporting Information
2S.
2.2 The Proposed Unifying Mechanistic Framework
2.2.1 Unifying Attributes: Pt&α-C, Pt&O, and Pt&H Interactions
The unifying mechanistic framework is predicated upon the use of attributes which
describe the interaction between an adsorption site and the anchoring atom of the
adsorbing molecule; and the dependence of these attributes on site geometry, structures
of adsorbed and dissolved species, and applied potential. These attributes are denoted by
Pt&α-C, Pt&O, and Pt&H interactions (“&” is used to indicate binding between Pt and
the adsorbed species through the α-C, O, or H atoms of the latter). Table 2.1 shows the
effects of site geometry (Fig. 2.1) and the type of adsorbed species formed at around 0.4V
(vs RHE). The basic premise is that species with very strong interactions (*CO and
*CRO) would block the sites extensively once they are formed. In their absence, species

Chapter 2
27
with comparable interactions can compete for adsorption, e.g. *COOH (Pt&α-C (others))
vs. *OCHO* (Pt&O) on Pt(111).
Table 2.1. Effects of Pt surface geometry on Pt&α-C, Pt&O, Pt&H interactions at ~ 0.4V.
Type of sites
Pt&α-C Pt&Oa
Pt&Ha
*CO b
, *CRO c
Others d
Pt(111) *T and Pt(111)-like *T Strong Moderate Moderate Weak
Pt(100) *T Very Strong Strong Moderate Moderate
Pt(110) *S and (110) *S on (111) Strong Moderate Strong Very Weak
Pt(110) *T and Pt(110)-like *T Very Strong Strong Very Weak Moderate
a. Supporting Information 2S1; b. Supporting Information 2S2; c. §2.2.4; d. §2.2.3 to
§2.2.5
Fig. 2.1.The surface geometry of Pt(100), Pt(111), Pt(110), and a plane with (110) steps
on (111) terraces (i.e. Pt(S)[(n-1)(111)x(110)], representing (n-1) rows of atoms on (111)
terraces before a (110) step. In this Fig, n = 3). Pt(110) is the plane with maximum (110)
step density on (111) terraces. *T on grey-colored atoms includes the Pt(100)*T, the
Pt(111)*T and the Pt(111)-like *T sites on Pt(S)[(n-1)(111)x(110)]. *T on orange-colored
atoms includes the Pt(110)*T and the Pt(110)-like *T sites on Pt(S)[(n-1)(111)x(110)].

Chapter 2
28
At potentials away from 0.4V, the relative strength of these interactions is still valid
column-wise but not row-wise. The significance of 0.4V is that it is close to the potential
of zero total charge (pztc) of Pt(111) and Pt(100) in 0.1M HClO4 [12, 13]. The adsorption
of *C-species at or around the pztc is the least affected by *H and *O-species
(Supporting Information 2S3). Below the pztc, Pt&H interaction is stronger than Pt&O
interaction and may be comparable to the Pt&α-C interaction other than those of strongly
bound *CO or *CRO. Above the pztc, the Pt&O interaction strengthens with increasing
potential, and becomes comparable to the strength of *CO at the onset potential of *OH
formation from water dissociation. The important potentials and the dominant species
from H2O which compete with *C-species for adsorption on different Pt sites in different
potential regimes are summarized in Table 2.2.
Table 2.2 The important potentials in 0.1M HClO4, and species from H2O dissociation
that compete with *C-species for adsorption.
Type of sites Low Potential High Potential
H2O*,OH2*
H*
pztca
H2O*
Onset of *OHb
HO*
Pt(111) *T ~0.37V 0.5~0.6V
Pt(100) *T ~0.42V ~0.6V
Pt(110)*S ~0.22V 0.3~0.4V
Pt(110)*T ~0.65V(pme)
a. pztc and pme (potential of maximum entropy, which is often slightly lower than pztc)
[12, 13](Supporting Information 2S1); b. onset of T*OH (i.e. onset of T*CO
electrooxidation) or onset of S*OH (i.e. onset of RCH2OH electrooxidation) (Supporting
Information 2S2).
2.2.2 CO Adsorption and Electrooxidation

Chapter 2
29
*CO is a common site blocker in the electrooxidation of oxygenates. It can only be
removed oxidatively by reaction with *OH to form *CO(*OH) and *COOH in sequence
(reaction 1). Due to the strength of T*CO adsorption (Tables 2.1-2.2), a high potential is
required to promote T*OH formation for it to be competitive with T*CO adsorption
and/or to react with the T*CO already formed. The onset of *CO electrooxidation on
Pt(111) and Pt(100) at ~0.5-0.6V therefore corresponds to the onset of T*OH formation
[12-14]. Since Pt&O interaction is stronger on the *S sites, *CO at low coverage would
not inhibit the step hollow sites for water activation [15]. Indeed S*OH may be able to
slowly oxidize adjacent T*CO at potentials below T*OH formation, albeit somewhat
slowly. This is evident from the presence of a small oxidation current in the low potential
region during CO stripping even though the main electrooxidation peak occurs around
0.8V [15-20] (Supporting Information 2S2).
↔ → (1)
2.2.3 HCOOH Adsorption and Electrooxidation
HCOOH electrooxidation at low potentials via the direct dehydrogenation pathway (§
2.2.3.1.) is the most desirable. However it only occurs when the catalyst surface is not
*CO inhibited. The goal in HCOOH oxidation is therefore to promote the direct
dehydrogenation pathway by suppressing *CO formation. This requires a good
understanding of how the various pathways are related to each other and influenced by
the Pt&α-C, Pt&O, and Pt&H interactions on different surfaces (Scheme 2.1).

Chapter 2
30
HCOOH
C
OHO
C
O
H
O
C
OOH C
O
O
H
C OO
C O
O H
O
H
H
solution
catalyzed
strong Pt&-C,
next to a sufficiently
strong Pt&O site
weaker Pt&-C
than Pt&O
sufficiently strong Pt&-C
No site blocking by *CO
optimal @ ~ pztc
For T*OH
stronger
than T*CO
direct pathway
Scheme 2.1.The proposed general reaction scheme for HCOOH electrooxidation. The
direct dehydrogenation pathway (CO2 formation via *COOH) is the most desirable for
current generation. It occurs when the surface is not blocked by *CO and is most
favorable when adsorption as *COOH is least interfered by H* and *O-species (i.e. at
around ptzc). T*CO formation can be minimized by a weaker Pt&α-C interaction; and by
the competing adsorption of species in the blue boxes. Once T*CO is formed, it can only
be removed effectively by oxidation when T*OH becomes abundant (i.e. at high V, via
the pathway in red).

Chapter 2
31
2.2.3.1 Dependence of Reaction Pathways on Pt&α-C, Pt&O, and Pt&H Interactions
In the absence of *CO inhibition, the adsorption of HCOOH as T*COOH requires a
sufficiently strong T*Pt&α-C interaction. Subsequent cleavage of the O-H bond in
T*COOH releases the hydrogen atom as H+
to water, with e-
passed to the electrode and
CO2 desorbed from the catalyst surface. This is known as the “direct dehydrogenation
pathway” of HCOOH electrooxidation [11].
The rate limiting step in the direct dehydrogenation pathway is the dehydrogenative
adsorption of HCOOH where the C-H bond in HCOOH is cleaved [21] to form *COOH.
This is more viable around the pztc when adsorption via the carbon atom is least affected
by H* and various *O-species (Supporting Information 2S3).
→ → (2)
If the Pt&α-C interaction is strong and there is an adjacent site with a sufficiently strong
Pt&O interaction (e.g. at moderate potentials), a bidentate transition state T*CO(*OH)
may be formed after dehydrogenative adsorption. A subsequent cleavage of the C-OH
bond leaves the surface with *OH and site-blocking T*CO. The *OH could also be
reduced and desorb as H2O if the potential is not sufficiently high to stabilize the *OH.

Chapter 2
32
→ → (3)
Although T*CO(*OH) forms when a strong T*Pt&α-C site is adjacent to sites with
sufficiently strong Pt&O interactions, it does not form when the adjacent sites are with
too strong a Pt&O interaction that stabilizes the *OH from water dissociation. Increasing
in potential increases the Pt&O interaction (Table 2.2). At potentials higher than the onset
of *OH formation from water dissociation, *OH becomes increasingly abundant and
reacts with T*CO to form T*COOH and CO2 in sequence (the reverse of the *CO
formation process). The greater presence of *OH also decreases the available *T sites and
adsorption as T*CO(*OH) (the precursor to T*CO formation) which requires two
contiguous sites. The increase in Pt&O interaction also increases the likelihood of
reversible adsorption as *OCHO*, which further diminishes the prospect of adsorption as
T*CO(*OH). By comparison the direct dehydrogenation pathway is not as adversely
affected by the competing adsorption from the *O-species since *COOH adsorption
requires only single site as opposed to T*CO(*OH) adsorption which requires dual sites.
As a consequence, the selectivity to the direct dehydrogenation pathway is enhanced
relatively in the presence of *O-species, thus preventing *CO inhibition and resulting in
an increase in current (Supporting Information 2S3).
On the other hand, *OCHO* is adsorbed with its C-H bond furthest away from the Pt
surface and is therefore more difficult to oxidize than *CO. Hence if T*CO is already
formed before Pt&O interaction is increased by raising the applied potential, the increase

Chapter 2
33
in Pt&O interaction allows HCOOH to adsorb as *OCHO* and competes with *OH
formation which is required for *CO removal. Hence T*CO electrooxidation is more
difficult in the presence of HCOOH than in an electrolyte without it (Supporting
Information 2S4).
-
(4)
2.2.3.2 Observations of Surface Geometry Dependency
Tables 2.1 and 2.2 above and the discussion in the preceding section may be used to
rationalize the following experimental observations:
1) Pt(110) has strong T*Pt&α-C and S*Pt&O sites next to each other and consequently the
T*CO(S*OH) adsorbed species can be established easily at low potentials. Recall that
T*CO(*OH) does not form when the Pt&O interaction is too strong that *OH from water
dissociation is stable, As a result, with increase in potentials that stabilizes S*OH, T*CO
can be formed via T*CO(T*OH) instead of T*CO(S*OH). Since the formation
(stabilization) of T*OH from H2O dissociation requires a high potentials (greater than
pme of ~0.65V in Table 2.2), T*CO formation via T*CO(T*OH) still proceeds at
moderately high potentials. Hence Pt(110) is the most active surface for T*CO formation,
as has been found in [22].

Chapter 2
34
2) Pt(100), with only moderate T*Pt&O interaction, does not favor the formation of
T*CO(T*OH) at low potentials. With a T*Pt&C interaction which is weaker than that on
Pt(110), the formation of T*CO(T*OH) is competed strongly by *O-species at high
potentials. As a result, Pt(100) is only active for T*CO formation around its pztc between
0.2V to 0.5V [22, 23].
3) Pt(111) with relatively moderate Pt&α-C and Pt&O interactions favors the adsorption
of HCOOH as *OCHO*, which is supported by the calculations in [24]. The adsorption
as T*CO(T*OH) is therefore inhibited. Furthermore, with a moderate Pt&α-C interaction,
desorption of *COOH as CO2 should be easier than on other planes. Pt(111) therefore has
the highest selectivity for the direct dehydrogenation pathway, as observed in [22].
2.2.4 Aldehyde Adsorption and Electrooxidation
A significant fraction of formaldehyde and acetaldehyde is hydrated in water and exists
in the diol forms of H2C(OH)2 and CH3CH(OH)2 [25]. The hydration and dehydration
reactions are reversible and immediately reach the equilibrium in water, especially water
with dissolved acidic and basic species [26]. This implies a fast hydration-dehydration
reactions between RCHO and RCH(OH)2 and low activation barrier for C-OH and O-H
bond cleavages from the >C(OH)2 structure in water. With this understanding and by
analogy with HCOOH electrooxidation, the mechanism for H2C(OH)2 electrooxidation
can be understood by means of Scheme 2.2.

Chapter 2
35
(5)
[H2C(OH)2] / [H2CO] = 2.28 x 103
, [CH3CH(OH)2] / [CH3CHO] = 1.1
at room temperature and pressure (r.t.p) [25, 27].
2.2.4.1 Major Difference between H2C(OH)2/ H2CO and HCOOH Electrooxidations
The adsorption and electrooxidation of H2C(OH)2/H2CO in Scheme 2.2 shares many
common features with HCOOH electrooxidation. There are, however, additional
pathways contributed by the dehydrogenative adsorption of H2CO (the dehydrated form)
to *CHO (and then *CO). Even without decomposing to *CO, *CHO has adsorption
strength comparable to that of *CO and is a site blocker except at high potentials where
*OH can add to it to form HCOOH (Supporting Information 2S5). As a result, catalyst
deactivation is more pervasive in H2C(OH)2/H2CO electrooxidation than in HCOOH
electrooxidation (as shown in [28]).
→ (6)
→ (7)
↔ (8)

Chapter 2
36
H2C(OH)2
C
OHHOC
O
H
O
H
H2CO
H
C
OHO
H
H
HCOOH
C
OHHO
C
OHO H
C OO
C
O H
C O
O
H
H
H
C
O
C
OHO
C O
O
H
H
C
O
O
H
weak Pt&C
no surface blocking
by *CO and *CHO.
strong Pt&C
adjacent site has
sufficiently strong Pt&O
direct pathway
high V
high V
Scheme 2.2. A proposed general reaction scheme for H2C(OH)2 electrooxidation. It is
analogous to HCOOH oxidation in the following aspects: direct dehydrogenation
pathways via O-H cleavage(s) in solution to form HCOOH and CO2, indirect pathways
via surface catalyzed C-OH cleavage forming inhibiting *CHO and subsequently *CO.
The main difference is the added possibility of *CHO formation from H2CO, which
makes surface inhibition an easier process.

Chapter 2
37
2.2.4.2 Similarities between H2C(OH)2 and HCOOH Electrooxidations
Similar to HCOOH, H2C(OH)2 electrooxidation also has direct dehydrogenation
pathways for current generation at low potentials, which occur through the
dehydrogenative adsorption of H2C(OH)2 as *CH(OH)2 when the surface is not
extensively blocked by *CO and *CHO. With subsequent O-H cleavage in solution and
surface-catalyzed C-H cleavage, HCOOH (reaction 9) and CO2 (reaction 10) are
eventually formed. Due to the ease of O-H cleavage in >C(OH)2 in water, the desorption
of *CH(OH)2 and :C(OH)2 to HCOOH and CO2 should be as viable as the desorption of
*COOH to CO2.
→ → (9)
-
→
- - -
→ (10)
Similar to *CO formation via the *CO(*OH) intermediate in HCOOH electrooxidation, a
strong Pt&α-C site next to a sufficiently strong Pt&O site can readily transform
*CH(OH)2 into *CH(OH)(*OH), followed by surface-catalyzed C-OH cleavage
to :CH(OH) and then O-H cleavage to *CHO in solution. In analogy to *CO(*OH) 
*CO with C=O converting to C≡O, *CH(OH)(*OH)  *CHO with C-OH converting to
C=O should also be feasible (See Supporting Information 2S6).
-
→
- - - -
→ (11)

Chapter 2
38
The similarities between H2C(OH)2 and HCOOH electrooxidations are reflected by their
similar surface geometry dependence: Pt(110) is the easiest to be deactivated by *CO,
followed by Pt(100) and Pt(111) in that order. The ranking is opposite to the viability of
the direct pathway on these surfaces [29, 30]. Furthermore, *OCH2O* has been detected
on Pt(111) at both low and high potentials (0.1V and 1.0V) [30]. The persistence of
*OCH2O* even at 1.0V indicates that it is as tenacious to oxidize as *OCHO*. Its
function on Pt(111) could be similar to that of *OCHO* - competing against adsorption
as *CH(OH)(*OH) and preventing the surface-catalyzed formation of *CHO and *CO at
low potentials.
2.2.4.3 Comparison between CH3CHO and H2CO
For acetaldehyde, the much lower aqueous phase equilibrium constant of
(~1.1) compared with (~2280) [25] suggests that there are more pristine
CH3CHO to dehydrogenate to *C(CH3)O than for H2CO to dehydrogenate to *CHO. Due
to the difficulty in *CRO oxidation, the surface is more readily deactivated in CH3CHO
than in H2CO (Supporting Information 2S7). Comparing the *CO formation from *CHO
and *C(CH3)O, the latter is more difficult because of the need to cleave the C-C bond of
*C(CH3)O to *CO and *CHx (reaction 12, most likely x=1 [31]). The cleavage of the C-C
bond is an arduous undertaking and requires sufficient free sites to bind to β-C at

Chapter 2
39
potentials below *OH formation [31-36]. The oxidation of *CO to CO2 and *CRO to
RCOOH (e.g. reaction 8) are categorized as “indirect pathways”.
→ → (12)
2.2.5 Alcohol Adsorption and Electrooxidation
2.2.5.1 The Pathways Determined by Pt&α-C and Pt&O Interactions
Scheme 2.3 summarizes the proposed mechanism for alcohol electrooxidation. Since
RCH2OH has only one O atom on its α-C, it requires reaction with *OH at potentials as
low as possible in order to form into CO2, RCOOH, or RCH(OH)2. The *S sites with
strong Pt&O interaction are the most suited to provide S*OH at lower potentials.
However, the Pt-C bond in the adsorbed species cannot be too strong, in order to allow
desorption soon after the formation of the second C-OH bond. We categorize these
reactions (reactions 13-14) as the “direct O-addition pathways” (More details in
Supporting Information 2S8). Reaction 13 is a hypothesis put forward for the first-time in
this thesis study.
→ → (13)
→ → → (14)

Chapter 2
40
RCH2OH
C
RHO
H
C
R
HO
H
O
H
RCH(OH)2
C
RHO
C
R
HO
O
H
C
OHHO
R
RCOOH
C
RO
C
ROH
C O
O
R
H
direct pathway:
weak Pt&-C
adjacent to *OH;
e.g. S*OH at low V
C
O
C
OHO
strong Pt&-C
R
C OO
H+
or
high V
high V
Scheme 2.3. Proposed reaction scheme for alcohol electrooxidation illustrating the direct
O-addition pathways to form carboxylic acid or hydrated aldehyde, and the formation of
inhibiting *CRO and *CO species. The presence of adjacent S*OH at low potentials and
an optimized Pt-C bond strength for desorption are required for high activity towards
direct O-addition pathways.

Chapter 2
41
A strong Pt-C bond does not only inhibit product desorption, but also the formation of the
second C-OH bond in reactions 13 and 14. This is because C-OH cleavage via
*CR(OH)(*OH) as in reaction 11 is favored by neighboring sites with very strong Pt&C
and sufficiently strong Pt&O interactions. The net result is the increased propensity
towards *CRO formation via reaction 15 and subsequently *CO formation via reaction 7.
→ (Supporting Information 2S6) (15)
Nevertheless, unlike RCHO that could adsorb directly as strongly bound *CRO, alcohol
is relatively weakly adsorbed and less favorable for *CRO formation [37, 38]. Hence,
alcohol electrooxidation is more active than aldehyde electrooxidation on Pt [28, 34].
2.2.5.2 Optimization of Surface Geometry and Operating Temperature
The (110) step density on (111) terraces should affect both the direct O-addition
pathways (onset potential of ~0.3-0.4V) and the formation of inhibiting *CO and *CRO
species. On these surfaces, the Pt(110)-like *T sites (the orange atoms forming *S sites
with the Pt layer below, Fig. 2.1) have very strong Pt&α-C interaction, and the Pt(111)-
like *T sites (Fig. 2.1) have relatively weaker Pt&α-C interaction. Therefore with a higher
(110) step density (i.e. more Pt(110)-like *T but fewer Pt(111)-like *T), the selectivity for
*CO and *CRO formation is increased (Supporting Information 2S9). For the increase of
current density at low potentials, more *S sites are needed to form S*OH for the direct O-
addition pathways, but these *S sites have to be kept away from the (110)-like *T in the

Chapter 2
42
same layer of the (111) terraces, in order to suppress *CRO and *CO formation. An
optimal step density therefore exists (Supporting Information 2S10). Nevertheless, if the
(110)-like *T sites can be avoided with the creation of *S such as those in the concave
surface shown in Fig. 2.2, the direct O-addition pathways may be optimized without the
side effect of *CRO and *CO formation.
Fig. 2.2 A concave surface with (111) terraces and (110) step hollow sites *S (red
triangles) but without the (110)-like *T sites.
However, for C2H5OH, if the selectivity towards C-C cleavage is the main concern, it
may be good to increase the (110) step density with more (110)-like *T, to catalyze the
formation of *C(CH3)O to promote the cleavage of its C-C bond into *CO and *CHx
(reaction 12). Furthermore, *C(CH3)O formation could also be enhanced by the
dehydration of CH3CH(OH)2 to CH3CHO (reverse of reaction 5) and probably the
*CCH3(OH)2 to *C(CH3)O at elevated temperatures, to suppress the direct O-addition
pathways forming CH3COOH and CH3CH(OH)2 (Supporting Information 2S11).
Together with the more facile kinetics of *CO electrooxidation at high temperatures [39],
improvements in both CO2 current efficiency and overall activity [40, 41] can be realized.

Chapter 2
43
2.3. Conclusion
A unifying framework for understanding the electrooxidation of formic acid, aldehydes,
and alcohols on Pt in acidic solutions has been proposed. Catalytic activity increases with
the resistance to C-OH bond cleavage on α-C bonded to two O atoms, or the ease of C-
OH formation on α-C bonded to one O atom. Pt&α-C, Pt&O and Pt&H interactions are
the most pertinent attributes to describe the effects of surface geometry (summarized in
Table 2.1) and applied potential (summarized in Table 2.2) on different reaction
pathways forming specific adsorbed species selectively.
For HCOOH which has two O atoms bonded to α-C, the main concepts for its
electrooxidation are summarized in Scheme 2.1 and are briefly described here: 1) the
direct dehydrogenation pathway via *COOH in the absence of *CO inhibition is the most
desirable for current generation, and is optimized when adsorption as *COOH is least
interfered by *H and *O-species; 2) surface blocking T*CO can easily be formed from
*CO(*OH) on Pt(110) and Pt(100) due to strong Pt&α-C and Pt&O interactions; 3) with
a weaker Pt&α-C interaction on Pt(111), and/or the presence of reversibly adsorbed
*OCHO* or/and *OH, T*CO(*OH) adsorption is hindered and so is the T*CO formation;
4) *OCHO* competes with *OH formation at high potentials and slows down the *CO
electrooxidation.

Chapter 2
44
Aldehyde (RCH(OH)2/RCHO) electrooxidation (summarized in Scheme 2.2) can also be
understood based on these basic concepts. The mechanism for the diol-form of the
aldehyde, RCH(OH)2 (with two O atoms) is largely similar to that of HCOOH
electrooxidation. However, the dehydrated form of aldehyde with one O atom (RCHO)
can adsorb as *CRO which is as surface inhibiting as *CO. *CRO can either form *CO at
low potentials or be oxidized by *OH to RCOOH at high potentials.
For alcohols with only one O atom attached to α-C, their electrooxidation is summarized
by Scheme 2.3. The addition of S*OH to T*CRHOH and :CROH on (111) terraces is
easier and hence RCH(OH)2 and RCOOH could be formed at lower potentials (onset
~0.3-0.4V), compared with the electrooxidation of strongly adsorbed *CRO and *CO
(onset ~0.6-0.7V). An optimum (110) step density on (111) terrace exists, since increased
step density supplies more S*OH at lower potentials but also creates more Pt(110)-like *T
with very strong Pt&α-C interaction that favors *CRO and *CO formation. Concave
surfaceswith (111) planes may provide the S*OH without Pt(110)-like *T. However, for
improving C-C cleavage and CO2 current efficiency concurrently in C2H5OH
electrooxidation, *C(CH3)O formation is desirable especially at elevated temperatures.

Chapter 2
45
CHAPTER 2 – SUPPORTING INFORMATION 2S
2S1. Pt&O and Pt&H (*H, *H2O, H2O*) Interactions at 0.4V
Pt&O interaction:
1) DFT calculations carried out in the absence of an electric field [42] indicate the
increase in the adsorption strength of H2O* in the following order: Pt(110) *T < Pt(111)
*T ~ Pt(100) *T < Pt(110) *S. This is an indirect indication of the increase in Pt&O
interaction in the same order.
Pt&H vs Pt&O interaction:
2) From laser-pulsed experiments in 0.1M HClO4 [12, 13], the potential at which
adsorbed H2O molecules orient as OH2* or H2O*, or the potential of maximum entropy
(pme) of double-layer formation, is 0.14V for Pt(110)*S, 0.33V for Pt(100) and 0.37V for
Pt(111). With increasing (110) step density on (111) terraces, the pme on the terraces
increases to about 0.65V. In addition, the potential of zero total charge (ptzc) below
which the presence of *H can be significant, is 0.22V for Pt(110)*S, 0.42V for Pt(100)
and 0.37V for Pt(111). These pme and pztc values are summarized in Table 2S.1 which

Chapter 2
46
also shows the dominant adsorbed species on Pt major crystallographic planes in 0.1M
HClO4.
Table 2S.1. The dominant adsorbed species on Pt basal planes in 0.1M HClO4
Pt sites on
basal planes
Increasing Potential
OH2*+H*
pme
H2O*+H*
pztc
H2O*
HO*
Pt(111) *T ~0.37V ~0.37V
Pt(100) *T ~0.33V ~0.42V
Pt(110)*S ~0.14V ~0.22V
Pt(110)*T ~0.65V
Dominant adsorbed species at 0.4V
Pt(111) *T H2O* Pt(110)*S H2O* + *OH
Pt(100) *T H2O* + H* Pt(110)*T *H2O + H*
A comparison between Pt&O and Pt&H interactions at 0.4V can then be made based on
Table 2S.1: With H2O* as the dominant adsorbed species on Pt(111) *T, Pt&O is slightly
stronger than Pt&H by comparison. The coexistence of H2O* and H* on Pt(100)*T
suggests comparable Pt&O and Pt&H interactions on this crystallographic plane. The
Pt(110)*S site is dominated by H2O* and *OH and hence Pt&O is much stronger than
Pt&H. On the other hand, *H2O and *H dominate on Pt(110)*T, indicating that Pt&H is
much stronger than Pt&O on this site.
The Pt&O and Pt&H interactions in Table 2.1 are ranked based on the understanding in
(1) and (2) above. Such interactions are to be compared with the Pt&α-C interaction to
determine the dominant adsorbed species under specific conditions. If adsorbed *C-

Chapter 2
47
species can be formed, Pt-&C must be at least comparable to or stronger than Pt&O and
Pt&H (2S-2).
2S2. Pt&α-C, Pt&O Interactions at 0.4V and around *OH Onset Potentials
1) The equilibrium surface coverage of *CO (θCO) depends on the adsorption potential (or
dosing potential, Ed). In Fig.2S.1 [17, 18], the high θCO values at 0.4V indicate that
Pt&CO interaction is much stronger than Pt&O and Pt&H. The higher θCO on Pt(100)
than on Pt(111) also indicates a stronger Pt&α-C interaction in the former.
2) From Fig. 2S.1, θCO generally decreases with the increase in adsorption potential.
When Ed is increased from 0.55V to 0.65V, θCO decreases sharply on Pt(111) and Pt(100)
[17, 18]. This potential range actually corresponds to the onset of T*OH formation
(reported to be 0.5V-0.6V for Pt(111) [12-14]) which competes with *CO adsorption and
oxidizes the adsorbed *CO.
3) For (110) steps on (111) terraces, S*OH onset may be around 0.3V (inference from the
low onset potential of RCOOH formation from RCH2OH, which requires reactions with
*OH [43, 44]). This indicates that the supply of S*OH at 0.4V should be quite plentiful.

Chapter 2
48
Fig.2S.1. Plot of CO-coverage on Pt(111) and Pt(100) surfaces in CO-free 0.1 M H2SO4
as a function of the dosing potential (squares). The total charge without double layer
correction (triangles), calculated from the hydrogen adsorption region of the
voltammogram, is also included [17, 18].
2S3. Suppression of *CO Formation and Optimization of the Direct *COOH
Pathway when Adsorption as *COOH is Least Affected by H* and *O-species.
1) For HCOOH electrooxidation on Pt(111) in the first forward voltammetric scan where
strongly bound *CO formation is insignificant, the observed current is limited by
adsorption as *COOH [21]. An oxidation peak develops around 0.4V during both
forward and reverse scan (Fig. 2S.2) [22], close to the pztc of Pt(111). The insignificant
hysteresis between the forward and reverse scans and the development of peak current
near surface pztc, are two characteristics of direct dehydrogenation pathways.

Chapter 2
49
Fig.2S.2. Cyclic voltammograms for two Pt basal planes in 0.1 M HCOOH + 0.1 M
HClO4. The solid lines represent first potential scans starting at 50 mV vs RHE. Dotted
lines correspond to the voltammogram in an electrolyte without HCOOH. Insets:
enlarged voltammograms in selected potential regions; units, mAcm-2
. Scan rate 50
mV/s.[22]
2) The reverse scan current (Fig. 2S.2) on Pt(100) is very high and a peak develops
around 0.42V (i.e. its pztc) after *CO is oxidized at high potentials [22, 23, 45]. When
potential decreases in the reverse scan, the leftover *OH could suppress the formation of
*CO(*OH) and *CO, thereby raising the selectivity for the direct dehydrogenation

Chapter 2
50
pathway via *COOH temporarily until most of the *OH desorbs at very low potentials.
With a strong Pt&C interaction to support effective dehydrogenative adsorption of
HCOOH (especially at its pztc), the turn-over rate on Pt(100) is therefore very high in a
short period of time.
3) Even though H2SO4 is known to cause specific anion adsorption which slows the
oxidation of *CO [46], a mixture of HClO4 and H2SO4 improves HCOOH
electrooxidation [47]. This could be explained by the sulfate inhibition of *CO(*OH)
adsorption and consequently *CO formation.
2S4. Observations of *OCHO* as an Inhibiting Species at High Potentials
1) *OCHO* has been implicated as a surface inhibiting species in an isotope study using
DCOOH & HCOOH[2], and by observations of the lowest measured methanol oxidation
current in the 0.7V-1.0V potential region concurrent with a high *OCHO* coverage on
Pt(111); relative to Pt(100) and Pt(110) which have higher *CO coverages [3]. The
strong adsorption of *OCHO* at high potentials and its inhibiting characteristic is similar
to the acetate adsorption (*OC(CH3)O*) [32, 48, 49].

Chapter 2
51
Fig.2S.3. Cyclic voltammogram for a 12
CO-covered Pt electrode in 0.5 M H2SO4 + 0.1 M
H13
COOH at a sweep rate of 50 mV/s; and the corresponding plot of the integrated band
intensities of *12
CO and *O13
CHO*in the positive-going scan (solid line). The dotted line
represents the oxidative removal of a *12
CO monolayer in an electrolyte without
H13
COOH [5].
2) Observations from reports which claim *OCHO* as an active intermediate in the
direct pathway [4-6] can in fact be interpreted as *OCHO* inhibition at high potentials.
For example Fig.2S.3 shows the voltammogram of a *12
CO-covered Pt electrode in
H13
COOH and a plot of the integrated band intensity in the forward scan of the
corresponding surface enhanced infrared adsorption spectrum (SEIRAS) [5]. At low
potentials the adsorption of H13
COOH is totally suppressed by the pre-adsorbed *CO
resulting in no measurable current. *CO oxidation is delayed to higher potentials in the

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