An unifying framework for understanding the electrooxidation of small organic molecules for fuel cell applications

... 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... 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. .. 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

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)

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

ACKNOWLEGEMENT I

1.2.2.1 Thermodynamic Cell Potential at Standard Conditions 4

1.2.2.3 Nernst Equation and Reversible Hydrogen Electrode (RHE) 6

1.2.3.1 Overpotential and Internal Resistance 7

1.4 The Capability of Proposed Unifying Mechanism and its Core Principles 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 2 MAJOR REACTION PATHWAYS IN THE

ELECTROOXIDATION OF SMALL OXYGENATES ON PLATINUM IN ACIDS

24

2.2.1 Unifying Attributes: Pt&α-C, Pt&O, and Pt&H Interactions 26

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

Electrooxidations 2.2.4.2 Similarities between H2C(OH)2 and HCOOH Electrooxidations 37 2.2.4.3 Comparison between CH3CHO and H2CO 38

2.2.5.1 The Pathways Determined by Pt&α-C and Pt&O Interactions 39 2.2.5.2 Optimization of Surface Geometry and Operating Temperature 41

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

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

2S12 Doubts in Recent Publications Supporting *OCHO* as Reactive

Intermediate

59

ACETALDEHYDE IN ACIDS AT HIGH POTENTIALS VIA ADSORBED CARBOXYLATES ON PLATINUM

3S1 Protracted *CO Electrooxidation in the Presence of Adsorbed Acetate 70

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

4.2.3 Effects of Pt/C Loading Per Electrode Surface Area 90

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

PATHWAYS IN ALCOHOL ELECTROOXIDATION ON BIMETALLIC PLATINUM-RUTHENIUM CATALYSTS

106

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

between 0.06V and 1.17V

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

6.2.1 Distribution of Sn and Its Effects on CO, Formaldehyde and Methanol

Electrooxidation

121

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

6.2.3.2 Inhibition by *OC(CH3)O* during Ethanol Electrooxidation 125

6S1 Temperature Effect on the Optimal Sn Distribution for Ethanol 130

CHAPTER 7 HIGH SELECTIVITY OF PALLADIUM CATALYSTS

FOR THE DIRECT DEHYDROGENATION PATHWAY IN FORMIC ACID ELECTROOXIDATION IN ACIDS

132

7.2.1 Strong Pd&H Interaction Results in Weak Pd&O Interaction 134

7.2.3 Enhanced Selectivity for the Direct HCOOH Pathway 138 7.2.4 Optimization of Pd-Based Catalysts for HCOOH Electrooxidation 141

ELECTROOXIDATION ON PLATINUM AND PALLADIUM IN ALKALINE SOLUTIONS

8.2.4.1 Optimization of Pt Catalysts 156

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

9.1 An Unifying Mechanistic Framework of Reactions 178

9.2 Considerations for Reactions in a Strongly Alkaline Environment 181

9.3.1 Comparison of Catalyst Activities 182 9.3.2 Important Indicators from Cyclic Voltammetry 182

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

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

*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

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

96

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

Table 1.1 Calculation of the ΔG0

and E0 for reactions 4 and 5 (ΔGf0: 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

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

Comparison of processes that affect the rate of ethanol

alkaline conditions on Pd

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

LIST OF FIGURES

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 12CO-covered Pt electrode in 0.5 M

H2SO4+ 0.1 M H13COOH at a sweep rate of 50 mV/s; and the

corresponding plot of the integrated band intensities of *12CO and

*O13CHO*in the positive-going scan (solid line) The dotted line

represents the oxidative removal of a *12CO monolayer in an

electrolyte without H13COOH

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

58

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

Fig 2S.7 Curve fitting for derived correlation “iformate α k (θ formate ) 3/2 / cHCOOH” 61

Fig 3.1 DEMS mass intensities of 12CO2 and 13CO2 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

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 electrooxidationof 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 HClO 4 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 HClO 4 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

CH CHO (A) or CH COOH (B) The catalyst loading was 5µg

127

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 /

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

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)

*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)

–H+–e

-a proton rele-ase from -adsorbed species vi-a inter-actions 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)

E0cell thermodynamic cell potential at standard conditions (V)

V applied or operating potential (V)

V cell operating cell potential (V)

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)

RHE Reversible Hydrogen Electrode

CE Coulombic Efficiency

EE Energy Efficiency

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

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 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

catalyst-fuel-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

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

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

2H2 + O2 2H2O (E0cell=1.229 – 0 = 1.229V) (3) where S.H.E is the acronym for the standard hydrogen electrode (vide infra), and E0cell 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

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

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

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

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

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

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

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

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

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

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

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