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

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) 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. I 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 II 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 Thesis Structure and Comparisons between Current and Proposed 15 1.5 Mechanisms Chapter 1 – Supporting Information 1S 1S1 Experimental CHAPTER 2 MAJOR REACTION PATHWAYS IN THE 21 21 24 ELECTROOXIDATION OF SMALL OXYGENATES ON PLATINUM IN ACIDS 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 31 Interactions 2.2.3.2 2.2.4 Observations of Surface Geometry Dependency Aldehyde Adsorption and Electrooxidation 33 34 III 2.2.4.1 Major Difference between H2C(OH)2/H2CO and HCOOH 35 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 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 Chapter 2 – Supporting Information 2S 43 45 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 48 *COOH Pathway when Adsorption as *COOH is Least Affected by H* and *O-species 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 55 Carboxylic Acids and Hydrated Aldehydes 2S9 Selectivity for *CO and *CRO Formation during Alcohol 57 Electrooxidation and Its Dependence on Step Density 2S10 Optimal (110) Step Density for Current Generation 58 IV 2S11 Elevated Temperature Enhanced Dehydration 59 2S12 Doubts in Recent Publications Supporting *OCHO* as Reactive 59 Intermediate CHAPTER 3 COMPLETE ELECTROOXIDATION OF ETHANOL AND 62 ACETALDEHYDE IN ACIDS AT HIGH POTENTIALS VIA ADSORBED CARBOXYLATES ON PLATINUM 3.1 Introduction 62 3.2 Proposed Mechanisms for the Complete Oxidation of Ethanol and 65 Acetaldehyde 3.3 Supporting Evidence for the Proposed Origin of the second CO2 Peak 67 3.4 Conclusion 68 Chapter 3 – Supporting Information 3S 70 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 75 *O*OCCO*O* CHAPTER 4 THE INHIBITION OF PLATINUM SURFACE BY 83 ACETALDEHYDE AND ACETIC ACID FORMATION DURING ETHANOL ELECTROOXIDATION IN ACIDS 4.1 Introduction 83 4.2 Results and Discussion 86 4.2.1 Electrooxidation of CH3COOH and CH3CHO 86 V 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 97 Tolerance 4.3 Conclusion Chapter 4 – Supporting Information 4S 98 100 4S1 CH3CHO electrooxidation at various concentrations 100 4S2 Effects of CH3CHO and CH3COOH Addition on Ethanol 102 Electrooxidation in Different Potential Regions 4S3 Observations that Support Direct O-Addition of Alcohol as the Major 104 Current Contributor in the Reverse Scan CHAPTER 5 PROMOTION OF THE DIRECT O-ADDITION 106 PATHWAYS IN ALCOHOL ELECTROOXIDATION ON BIMETALLIC PLATINUM-RUTHENIUM CATALYSTS 5.1 Introduction 106 5.2 Results and Discussion 108 5.2.1 Observations Supporting the Enhancement of the Direct O-addition 108 Pathways 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 VI 5.2.5 PtRu Activation by Cyclic Voltammetric Pretreatment in C2H5OH 115 between 0.06V and 1.17V 5.3 Conclusion 118 Chapter 5 – Supporting Information 5S 5S1 The Activation of Deactivation 119 of PtRu during Methanol 119 Electrooxidation CHAPTER 6 EFFECTS OF TIN IN PLATINUM-TIN CATALAYSTS 120 FOR ELECTROOXIDATION IN ACIDS 6.1 Introduction 120 6.2 Review, Results and Discussion 121 6.2.1 Distribution of Sn and Its Effects on CO, Formaldehyde and Methanol 121 Electrooxidation 6.2.2 Similarity between Methanol and Ethanol 6.2.2.1 123 Enhancement of the O-Addition Pathway for Alcohols by *OH on 123 Sn/SnOx and Weaker Pt&α-C Interaction 6.2.2.2 6.2.3 Adsorption is Rate-limiting on Pt-Sn Alloys Difference Between Ethanol and Methanol 124 125 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 6.3 Comparison between Pt-Sn and Pt-Ru Conclusion Chapter 6 – Supporting Information 6S 6S1 127 128 130 Temperature Effect on the Optimal Sn Distribution for Ethanol 130 VII Electrooxidation CHAPTER 7 HIGH SELECTIVITY OF PALLADIUM CATALYSTS 132 FOR THE DIRECT DEHYDROGENATION PATHWAY IN FORMIC ACID ELECTROOXIDATION IN ACIDS 7.1 Introduction 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 CHAPTER 8 EFFECTS OF IONIZATION ON ETHANOL 143 144 ELECTROOXIDATION ON PLATINUM AND PALLADIUM IN ALKALINE SOLUTIONS 8.1 Introduction 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 153 Pathways 8.2.2 Electrooxidation of CO 154 8.2.3 C1 Pathways on Pt and Other Poisoning Species 154 VIII 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 CHAPTER 8 – Supporting Information 8S 8S1 157 159 Weaker Effects of Acetic Acid on Pt catalysis under Strongly Alkaline 159 Conditions 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 163 Solutions 8S3 Formation of CH3CHO by the Direct Dehydrogenation of CH3CH2O- 8S4 Effects of Acetaldehyde and pH on Pd Catalysis in Strongly Alkaline 166 165 Solutions 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 9.1 CONCLUSIONS AND RECOMMENDATIONS An Unifying Mechanistic Framework of Reactions 177 178 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 IX 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 REFERENCE Recommendations 183 186 X 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 XI 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 XII 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. XIII LIST OF SCHEMES Scheme 2.1 The proposed general reaction scheme for HCOOH 30 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). Scheme 2.2 A proposed general reaction scheme for H2C(OH)2 electrooxidation. 36 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. Scheme 2.3 Proposed reaction scheme for alcohol electrooxidation illustrating 40 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. Scheme 3.1 The proposed pathways (non-elementary steps) for the complete 66 oxidation of C2H5OH and CH3CHO to CO2 in different potential regions. Scheme 4.1 Suggested reaction scheme for C2H5OH electrooxidation. R is CH3. 96 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 XIV 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 113 (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). Scheme 8.1 The reaction mechanism from reference for ethanol 145 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. Scheme 8.2 Proposed reaction mechanism for ethanol electrooxidation on Pt. 149 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. XV LIST OF TABLES 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 27 at ~ 0.4V. Table 2.2 The important potentials in 0.1M HClO4, and species from H2O 28 dissociation that compete with *C-species for adsorption. Table 2S.1 The dominant adsorbed species on Pt basal planes in 0.1M HClO4 Table 3.1 Summary of the ethanol reaction mechanisms showing the effects of 68 Pt&O, Pt&α-C, Pt&β-C interactions on various electrooxidation pathways. 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 116 Electrooxidation and Reduction of Surface Species in 0.1M HClO4 Table 6.1 Distribution of Sn/SnOx among the Pt atoms on the catalyst surface 122 and their effects on the electrooxidation of CO, aldehydes, and alcohols* Table 6S.1 Product distribution in the effluent of single cell tests at 90˚C Table 8.1 Comparison of processes that affect the rate of ethanol 148 electrooxidation on Pt under acidic and strongly alkaline conditions 46 131 XVI Table 8.2 Comparison of ethanol electrooxidation under acidic and strongly 151 alkaline conditions on Pd Table 8S.1 Effects of CH3CHO addition to ethanol electrooxidation at pH 13.93 171 and 13.40 Table 9.1 Summary of Pt-catalyzed electrooxidation of different oxygenates 179 in acidic solutions at room temperature Table 9.2 Summary of the effects of different Pt-based catalysts and operating 180 conditions XVII 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)[(n1)(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 48 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. Fig. 2S.2 Cyclic voltammograms for two Pt basal planes in 0.1 M HCOOH + 49 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. 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. Fig. 2S.4 Potential oscillations observed in 0.5 M H2SO4+ 0.1 M formaldehyde 53 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. Fig. 2S.5 CVs of Pt single crystals in 0.5 M CH3OH and 0.5 M HClO4 at a scan 58 rate of 2mV/s: (a) Pt basal planes, (b) Pt surfaces with (110) steps on (111) terraces. 51 XVIII 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 12CO2 and 13CO2 during oxidative stripping 64 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). Fig 3S.1. (a) Integrated IR intensities of *CO from pre-adsorbed CO and 71 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. 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) preadsorbed 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 88 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. 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 XIX 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 95 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 ) Fig. 4S.1 Cyclic voltammograms of a E-TEK Pt/C catalyst with 20µg Pt /cm2 101 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. Fig. 5.1 (A) Steady-state cyclic voltammograms of electrooxidation of ethanol 110 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 Fig. 5.2 Cyclic voltammogram of 1M ethanol electrooxidation in 0.1M HClO4 on 5µg 112 PtRu /cm2 with an anodic scan limit of 0.7V. Scan rate: 10mV/s Fig. 5.3 Cyclic voltammograms in 0.1M HClO4 on 20µg PtRu and 20µg Pt 115 /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 Fig. 5S.1 Cyclic voltammogram of 1M methanol electrooxidation in 0.1M HClO4 on 119 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. Fig. 6.1 The 20th scan cyclic voltammograms of the electrooxidation of 1M 127 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 XX 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. (– 136 ) 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. Fig. 7.2 The effect of adsorption potential of CO on charge passed to the 137 electrode during CO adsorption (Qads) and during subsequent electrooxidation of *CO (QOx). Fig. 7.3 Voltammograms of the electrooxidation of formic acid on Pd(111) and 140 Pd(100) in 0.1 M HClO4 containing 0.1 M formic acid. Scanning rate: 20 mV/s Fig. 7.4 Voltammograms of formic acid electrooxidation on modified Pd 142 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). Fig. 8.1 Stabilized cyclic voltammograms of ethanol electrooxidation in 0.1M 147 HClO4 (with/without rotation at 1000rpm) and in 0.85M KOH. Catalyst loading: 5µg Pt/cm2. Scan rate: 10mV/s. Fig. 8.2 Effect of acetic acid (A) and acetaldehyde (B) addition on stabilized 150 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. Fig. 8.3 Steady state voltammograms of ethanol electrooxidation in 0.85M 153 KOH electrolyte with different CH3COO- concentrations. Catalyst loading: 5µg Pd/cm2. Scan rate: 10mV/s. Fig. 8S.1 Effect of pH on the voltammogram of Pt(111) at 30 mV/s in (a) 20 160 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)). Fig. 8S.2 Voltammograms of ethanol oxidation on E-TECK Pt/C (20µg Pt / 161 cm2) in 0.85M KOH with the addition of different amounts of XXI CH3COOH which leads to 0.85M K+ and final CH3COOconcentrations 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 163 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. Fig. 8S.4 Effect of CH3CHO addition to stabilized voltammograms of C2H5OH 164 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. Fig. 8S.5 Linear sweep voltammograms of (a) alcohols (10mM) with high j, (b) 166 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. Fig. 8S.6 Effect of acetaldehyde addition on the voltammogram of E-TEK Pd/C 168 (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. Fig. 8S.7 Effects of acetaldehyde addition on the voltammogram of E-TEK 170 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. Fig. 8S.8 Effect of E-TEK Pd/C loading per electrode surface on the 172 voltammogram of 0.85M KOH + 1M ethanol at 10mV/s. Current density is normalized by a) Pd mass, or b) electrode geometrical area. Fig. 8S.9 Effect of E-TEK Pt/C loading per electrode surface on the 173 voltammogram of 0.85M KOH + 1M ethanol at 10mV/s in. Current density is normalized by a) Pd mass, or b) electrode geometrical area. Fig.8S.10 Voltammograms of ethanol electrooxidation on Pt(111) and Pt(110) 175 for the 1st (a) and the 20th (b) cycle in 0.5M ethanol and 0.1M NaOH, at 10mV/s. XXII 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) E0cell thermodynamic cell potential at standard conditions (V) XXIII 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) ) XXIV SHE Standard Hydrogen Electrode RHE Reversible Hydrogen Electrode CE Coulombic Efficiency EE Energy Efficiency XXV Chapter 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 ondemand 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. 1 Chapter 1 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-fueloperating 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. 2 Chapter 1 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. 3 Chapter 1 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 (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 4 Chapter 1 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. (ΔGf0: Standard Gibbs free energy of formation of compounds, from [1]) Compound ΔGf0 (kJ/mol) ΔG04 (kJ/mol) E04 CO2 (g) H2O (l) C2H5OH (l) -394.4 -237.1 -174.8 (-174.8) + 3(-237.1) – 2(-394.4) = -97.3 -(-97.3)(1000)/12/96485 = 0.084V ΔG05 (kJ/mol) 2(-394.4) + 3(-237.1) – (-174.8) = -1325.3 0 -(-1325.3)(1000)/12/96485 = 1.145V Alternatively, E01 (=1.229) - E04 = 1.145V E 5 5 Chapter 1 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 6 Chapter 1 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 7 Chapter 1 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. 8 Chapter 1 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). 9 Chapter 1 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. 10 Chapter 1 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 11 Chapter 1 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 fuelcatalyst-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 12 Chapter 1 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. II. interactions between the catalytic site and adsorbed *H, *C-species and *O-species 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, *Cspecies 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. 13 Chapter 1 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 14 Chapter 1 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 15 Chapter 1 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 16 Chapter 1 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. 17 Chapter 1 Table 1.3. Comparison between Current and Proposed Mechanisms Chapter 2 (Reaction Pathways on Pt at Practical Anode Potentials) Fuel Molecule Current Mechanisms Reactive Intermediate Blocking Intermediate *COOH *CO *OCHO* HCOOH *COOH, H2CO / H2C(OH)2 *OCH2O* CH3CHO / CH3CH(OH)2 Non CH3OH *CO *CO, *C(CH3)O (*OCHO* only weakly suppresses the *COOH pathway) Desorbed Product T*CO T*CH(OH)2 T*CO, :C(OH)2 T*CHO Non T*CO, *C(CH T 3)O CO2 HCOOH CO2 Non HCOOH T*CO, *CO *C(CH3)O + *OH  CH3COOH CH3CH2OH Blocking Intermediate :CHOH + S*OH  T*CH(OH)2 CH3O* – H*H2CO :C(CH3)OH + *OH  *C(CH3)(OH)2 Reactive Intermediate T*COOH (*OCHO* is not *CO reactive but suppresses *CO *CO, *OCHO* formation) :CHOH + *OH  *CH(OH)2 *CH2OH (unknown pathway) Proposed Mechanisms T*CH2OH + S*OH  H2C(OH)2 T*CHO :C(CH3)OH + S*OH  T*C(CH3)(OH)2 *CO H2C(OH)2 CH3COOH T*CO, T*C(CH3)O CH3CH2O* – H*  CH3CHO T*CH(CH3)OH CH3CH(OH)2 + S*OH  CH3CH(OH)2 *CH(CH3)OH – H*  CH3CHO Chapter 3 (Reaction Pathways at Potentials Higher than *CO Electrooxidation on Pt) Fuel Molecule CH3CH2OH, CH3CHO Current Mechanisms Intermediate at High Potentials *CHx Conflicting Findings Proposed Mechanisms Intermediate at High Potentials Remark They suppress *CHx was found *O-species, e.g. *OH formation, but their to convert to *OCHO*, oxidation *CO at low *O*OCCO*O*, require *OH potentials *C(CH3)O* Desorbed Product CO2 18 Chapter 1 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; *C(CH T 3)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 PtCO 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+ + eT*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 Reactive Pathway *C(CH3)O + *OH  CH3COOH CH3CH2OH / CH3CH2O- Proposed Mechanisms (for Pt) Reactive Pathway :C(CH3)O- + S*OH  T*C(CH3)(OH)O (negligible surface inhibition by *C(CH3)O) For Pd, with to stronger Pd&H but relatively weaker Pd&C: :C(CH3)O- + *OH-  T*C(CH3)(OH)O + e Desorbed Product CH3COO- CH3COO- 19 Chapter 1 CH3CH2O* – H*  CH3CHO *CH(CH3)OH – H*  CH3CHO CH3CH(OH)O/ CH3CH(OH)2 / CH3CHO / CH2CHO- - + S*OH  CH3CH(OH)O- T*CH(CH3)O T*CH(CH3)O -  CH3CHO + e- CH3CH(OH)OCH3CHO *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. 20 Chapter 1 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 21 Chapter 1 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 22 Chapter 1 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. 23 Chapter 2 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 24 Chapter 2 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. 25 Chapter 2 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 26 Chapter 2 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(111) *T and Pt(111)-like *T Pt(100) *T Pt(110) *S and (110) *S on (111) Pt(110) *T and Pt(110)-like *T Pt&α-C *CO , *CRO c Strong Very Strong Strong Very Strong b Pt&Oa Pt&Ha Moderate Moderate Strong Very Weak Weak Moderate Very Weak Moderate d Others Moderate Strong Moderate Strong 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)]. 27 Chapter 2 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 High Potential Onset of *OHb 0.5~0.6V ~0.6V 0.3~0.4V HO* H2O* H2O*, OH2* H* Pt(111) *T Pt(100) *T Pt(110)*S Pt(110)*T Low Potential pztca ~0.37V ~0.42V ~0.22V ~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 28 Chapter 2 *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). 29 Chapter 2 HCOOH weaker Pt&-C than Pt&O sufficiently strong Pt&-C No site blocking by *CO optimal @ ~ pztc HO H C C O O strong Pt&-C, next to a sufficiently strong Pt&O site direct pathway solution catalyzed H O O H H O O O O C C H For T*OH stronger than T*CO O C O H C O O 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 Chapter 2 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. 31 Chapter 2 → → (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 32 Chapter 2 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) T*CO(*OH) adsorbed species can be established easily at low potentials. Recall that 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]. 33 Chapter 2 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. 34 Chapter 2 (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) 35 Chapter 2 H2C(OH)2 H2CO weak Pt&C no surface blocking by *CO and *CHO. H H H C O HO OH C O strong Pt&C adjacent site has suf f iciently strong Pt&O direct pathway HO H OH H C HO O H OH H C O H O C O H H H O O C HCOOH C H C O high V HO O H C O C O O C O high V C O 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 Chapter 2 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) 37 Chapter 2 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 38 Chapter 2 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) 39 Chapter 2 RCH2OH strong Pt&-C H HO R HO C R H O R C C direct pathway: weak Pt&-C adjacent to *OH; e.g. S*OH at low V H HO C R R H HO O C H O O R RCH(OH)2 HO C R C OH high V R O H H+ or O C C O R RCOOH high V HO O C O C O 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 Chapter 2 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 Oaddition pathways, but these *S sites have to be kept away from the (110)-like *T in the 41 Chapter 2 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. 42 Chapter 2 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 COH 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. 43 Chapter 2 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. 44 Chapter 2 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 45 Chapter 2 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 ~0.33V ~0.14V ~0.37V ~0.42V HO* Pt(110)*T ~0.37V H2O* Pt(110)*S pztc H2O* + H* Pt(100) *T pme OH2* + H* Pt(111) *T Increasing Potential ~0.22V ~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 *C46 Chapter 2 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. 47 Chapter 2 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. 48 Chapter 2 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 49 Chapter 2 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]. 50 Chapter 2 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 [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 *12CO-covered Pt electrode in H13COOH 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 H13COOH is totally suppressed by the pre-adsorbed *CO resulting in no measurable current. *CO oxidation is delayed to higher potentials in the 51 Chapter 2 presence of HCOOH. Since *CO oxidation requires the presence of *OH, the delay is an indication of the difficulty in *OH generation caused by an inhibiting species originated from HCOOH. Since the isotope study did not detect *13CO formation from H13COOH, inhibiting species competing with *OH adsorption has to be *O13CHO*. 3) Fig. 2S.4 [6] shows the potential oscillations during the electrooxidation of formaldehyde under galvanostatic conditions and the surface coverage by various reaction intermediates from SEIRAS measurements. The potential spike correlates well with high *OCHO* and low *CO coverage and is an indication that *OCHO* is more difficult to oxidize (i.e. more inhibiting) than *CO, thereby requiring higher potentials for its removal. Similar oscillations have also been observed in formic acid electrooxidation [4]. If *OCHO* were an active species at low potentials as claimed in [6], then a high *OCHO* coverage and low *CO coverage should give rise to a facile reaction, and not a potential spike that indicates kinetic difficulty. We hypothesize that *CO electrooxidation and *OCHO* formation occur during the rising transient. On the other hand, during the falling transient, *OCHO* removal by oxidation instead of *CO removal is favorable. The different events occurring on opposing sides of the potential spike have to come from the influence of another adsorbed species, probably *OH. The experimental results are consistent with *OH accumulation and *CO oxidation in the rising transient; and the oxidation of *OCHO* after *OH reaches sufficient coverage at the spike potential. Once *OCHO* oxidation begins, free sites are regenerated and direct dehydrogenation pathways can then proceed to lower the potential. The reduction in potential consequently desorbs *OCHO* and *OH so that *CO accumulation can recommence instantaneously. 52 Chapter 2 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 [6]. 2S5. *CHO as one of Surface Blocking Species Despite the onset potential of H2C(OH)2 oxidation (~ 0.2V) via direct pathway on Pt(111) is quite low, it only appeared in the first cycle of cyclic voltammogram since the current density dropped tremendously from second cycle especially at 0.6V and below [30]. The surface blocking *CO was found to be the cause of this deactivation, but there was no evidence to exclude *CHO as another potential surface blocking species. We suggest 53 Chapter 2 *CHO to be surface blocking due to its similar structure and property to *C(CH3)O, which is known to adsorb strongly on Pt until it starts to be oxidized at high onset potential around 0.7V [32, 33, 50]. Besides, the potential spike in Fig. 2S.4 [6] corresponds well with the quick accumulation of *OCHO* that begin at 0.6V. This potential is much higher than the onset of HCOOH formation (~ 0.2V) via direct pathway on Pt(111) [30], suggesting that HCOOH formation with immediate re-adsorption as *OCHO* at potentials >0.6V may come from *CHO electrooxidation. Besides, the 0.6V in Fig. 2S.4 also corresponds to acute *CO electrooxidation (sharp drop in *CO). Since 0.6V is around the potential for T*OH formation, both *CO and *CHO likely require T*OH for their electrooxidation and they could be equally surface blocking. 2S6. Conversion of :CROH to *CRO Both :C(CH3)OH from C2H5OH [51] and :CHOH from CH3OH [52] have been detected on Pt. :C(CH3)OH is sensitive to dissolved oxygen and is easily oxidized to *C(CH3)O [51]. 2S7. Stronger Surface Inhibition by CH3CHO than by H2CO The electrooxidation of acetaldehyde provides a good example for RCOOH formation via *CRO. Even for the direct-dehydrogenation-pathway efficient Pt(111), acetic acid is only 54 Chapter 2 formed from acetaldehyde at an onset potential around 0.7V, which is significantly higher than the potential for the direct H2C(OH)2 electrooxidation to HCOOH (onset~0.2V) [30, 32, 33, 50] as well as our suggested onset potential for electrooxidizing *CHO (~0.6V in §2S5). This suggests the absence of a significant direct dehydrogenation pathway for CH3CH(OH)2 electrooxidation due to the lower ratio than ratio and a stronger inhibition by *C(CH3)O than *CHO on the catalyst surface. 2S8. Direct O-Addition Pathways in the Oxidation of Alcohols to Carboxylic Acids and Hydrated Aldehydes Direct O-addition pathways in alcohol electrooxidation produce carboxylic acids and hydrated aldehydes by reacting *OH with :CROH or with *CRHOH (reactions 13-14 in Chapter 2 main text), without formation of *CRO nor *CO as adsorbed intermediate. 1) Quite a number of C2H5OH electrooxidation studies reported formation of CH3COOH at potentials < 0.6V [44, 49, 53-57]. For Pt(111) and Pt/C nanoparticles the onset can be as low as 0.3-0.4V [44, 49, 55]. However, the onset potential for CH3COOH formation from acetaldehyde via *C(CH3)O is around 0.7V [32, 33, 50]. The onset potential for C2H5OH oxidation to acetaldehyde also tends to be slightly higher than the onset for C2H5OH oxidation to CH3COOH [44, 53]. These are indications of the existence of other 55 Chapter 2 pathways which forms CH3COOH from C2H5OH directly without *C(CH3)O as the reaction intermediate. 2) There are significant inconsistencies on the values of onset potentials for Pt(111). For CH3OH oxidation in HClO4, it has been reported as 0.55V or 0.35V [52, 58], although both publications reported a consistent onset potential for Pt (110) ~ 0.45V, thus eliminating experimental errors in potential measurements as the cause for error. For C2H5OH oxidation, the onset of CH3COOH formation was reported to be around 0.3V [44, 49] or 0.6V on Pt(111) [59]. This may be explained by the recent discovery that the C2H5OH electrooxidation activity of Pt(111) is very low in the first cycle before defect creation [57]. Surface with low onset potentials are most likely non-ideal Pt(111) with surface defects or (110) steps. The proposed CH3COOH formation mechanism at low potentials is the same as our reaction 14 (Chapter 2). Similar mechanism for HCOOH formation from CH3OH was also suggested based on recent calculations performed for Pt(111) [60]. However that study was unable to differentiate the difficulty in oxidizing *CHO and :CHOH. 3) Since RCHO forms strongly adsorbed *CRO easily, during RCH2OH electrooxidation it is unlikely for any adsorbed intermediate to directly desorb into RCHO. The following is a list of observations supporting RCH(OH)2 as the desorbed product via reaction 13 (Chapter 2), a direct O-addition pathway similar to CH3COOH formation: (i) among the Pt major crystallographic planes, acetaldehyde and CH3COOH are formed most easily on 56 Chapter 2 Pt(111) (in our opinion, Pt(111) with defects) at low potentials [49]; (ii) the acetaldehyde yield improves at higher C2H5OH concentrations but an optimal concentration exists for the CH3COOH yield, indicating acetaldehyde and CH3COOH formations are parallel reactions with acetaldehyde requiring less clean Pt sites [44, 53]; (iii) in a study using *CO free but *OH rich surfaces (created by stepping the potential from high to low values), indicates that the rate limiting step in CH3OH electrooxidation is C-H cleavage into *CH2OH for Pt(111) and Pt(110), and C-H cleavage into :CHOH for Pt(100)[61]. These are the intermediates in our proposed pathways (reactions 13-14, Chapter 2) before the fast addition of another OH bond promoted by this *OH rich surface. 2S9. Selectivity for *CO and *CRO Formation during Alcohol Electrooxidation and Its Dependence on Step Density 1) At high potentials, CH3OH electrooxidation leads to a higher *CO coverage on Pt(110) than on Pt(100); which is significantly higher than that on Pt(111)[3, 52]. The situation for C2H5OH electrooxidation is similar: Pt(110) is active for C-C bond cleavage and *CO formation over a wider potential range (0.1V – 0.8V); Pt(100) is only active for *CO formation at low potentials up to 0.6V, which is the onset for T*OH formation for *CO and *C(CH3)O electrooxidation into CO2 and CH3COOH; Pt(111) is very poor at C-C cleavage [49]. 2) For CH3OH electrooxidation, the peak potential increases with the (110) step density as Pt(111) < Pt(554) < Pt(553) < Pt(110) (Fig. 2S.5) [58]. Similarly for the C2H5OH 57 Chapter 2 electrooxidation, the peak potential increases as Pt(111) ~ Pt(554) < Pt(553) < Pt(110) [34, 62]. The increase in peak potential actually indicates a stronger Pt-C bond in an oxidizable adsorbed species, and the delay in inactive surface oxide formation to higher potentials. This implies that more alcohol is adsorbed as *CRO and *CO and oxidization has to occur at higher potentials. 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]. 2S10. Optimal (110) Step Density for Current Generation The current density for CH3OH electrooxidation increases in the order of Pt(111) < Pt(110) < Pt(554) < Pt(553) in both HClO4 (Fig.2S.5) and H2SO4 electrolytes [58], indicating an optimal (110) step density. For the C2H5OH electrooxidation, there is also an optimal (110) step density for maximizing peak current density in HClO4, i.e. Pt (554) [34, 62]. 58 Chapter 2 2S11. Elevated Temperature Enhanced Dehydration It is known that the equilibrium constant of [CH3CH(OH)2] / [CH3CHO] decreases with increasing temperature [27]. Besides, elevated temperature also enhances surfacecatalyzed dehydration of HCOOH as *CO [63]. 2S12. Doubts in Recent Publications Supporting *OCHO* as Reactive Intermediate In recent studies on electrooxidation mechanisms with in-situ experiments such as SEIRAS to determine the surface component during voltammetry, correlation between *OCHO* signal strength (which is proportional to its surface coverage) and the measured current density is often derived to show *OCHO* to be the reactive intermediate. However, due to the complex reaction dynamics, both the measured current density and the *OCHO* signal strength may be affected by many reasons, such as the presence of *CO, or/and *OH. It is hence difficult to conclude a reaction mechanism based on correlation (as in 2S4 we explained the Fig. 2S.4 in a different way from its original paper). Osawa et. al. tried to avoid the effect of *CO, and hence tried to correlate the measured current to the *OCHO* signal strength at 0.5s after the potential step down from 0.9V (to oxidize *CO) to 0.6V, as shown in Fig. 2S.6 and Fig. 2S.7 [10]. However, a careful examination on the rate equation they derived for the correlation, there are two doubtful assumptions: i) there is insignificant *OH coverage at 0.6V, which is unlikely at just 0.5s after stepping down from 0.9V (where a high *OH coverage is needed to oxidize *CO); ii) 59 Chapter 2 steady state when formate coverage does not change over time, however it can be seen from Fig. 2S.6 that the current is still decreasing at 0.5s and the high fluctuation in formate signal is difficult to confirm the absence of minor change in formate coverage. 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 [10]. Furthermore, they tried to correlate their derived equation “iformate α k (θformate)3/2 / cHCOOH ”, as “iformate α k (Aformate)3/2 / cHCOOH” in Fig. 2S.7, with only four points on the straight line. Although it was explained that at low HCOOH concentration (cHCOOH), their derived correlation is invalid due to competition with (bi)sulfate adsorption, it is in fact mathematically questionable to simply replace the formate coverage (θformate) by the formate signal (Aformate), as there is a power of 1.5 to it. 60 Chapter 2 Fig. 2S.7 Curve fitting for derived correlation “iformate α k (θformate)3/2 / cHCOOH ” [10]. The debate on *OCHO* or *COOH as reactive intermediate may still be going on. However, we would like to invite the reader to analyze whether the *OCHO* as reactive intermediate is able to explain the below phenomena: i) Why is backward scan of Pt(100) much higher than that of Pt(111)? (we have explained in 2S3-2) ii) During potential oscillations (Fig 2S.4), what is the contribution to the potential spike if it is not due to difficult oxidation of *OCHO*? iii) Why is Pd much more active for HCOOH electrooxidation than Pt? (we will discuss in Chapter 7) 61 Chapter 3 CHAPTER 3 COMPLETE ELECTROOXIDATION OF ETHANOL AND ACETALDEHYDE IN ACIDS AT HIGH POTENTIALS VIA ADSORBED CARBOXYLATES ON PLATINUM 3.1. Introduction Ethanol is an attractive renewable energy carrier for fuel cells. However the activities of current catalysts for ethanol electrooxidation are unsatisfactory. The major reaction products with most catalysts are acetaldehyde and acetic acid rather than CO2, and hence the number of electrons that can be extracted from the ethanol oxidation reaction (EOR) is much lower than the theoretical maximum (12). Since Pt is the most common catalytic metal for electrooxidation in acids, it is important to have a lucid understanding of the EOR on Pt before we can make use of adjuvant components (e.g. Ru, Rh, and metal oxides [64-66]) effectively in the catalyst design. Contrary to the low selectivity for C-C bond cleavage in the electrochemical oxidation of bulk C2H5OH, CO2 has always been the only desorbed product detected by differential electrochemical mass spectrometry (DEMS) during oxidative stripping voltammetry of adsorbed residues from ethanol [31, 32, 38, 51, 67, 68], acetaldehyde [32, 33, 35, 36, 38], propanol isomers [69-71], butanol isomers [72], ethylene glycol and its C2 derivatives 62 Chapter 3 (glycol aldehyde, glyoxal, glycolic acid, glyoxylic acid, oxalic acid) [37]. In these investigations, the test molecule was adsorbed for several minutes at a predetermined potential, i.e. the adsorption potential Ead. The electrolyte was then exchanged with a fresh electrolyte (HClO4 or H2SO4) at the same applied potential. There was no indication of irreversible poisoning of the catalyst surface since the surface after oxidative stripping could recreate the voltammetric response of a clean surface. There are generally two CO2 peaks in the stripping voltammogram; as shown in Fig. 3.1 for the oxidative stripping of isotopically labeled ethanol and acetaldehyde. DEMS was used to identify the origin of CO2. The first CO2 peak at lower potential (~0.7V) has been quite extensively investigated by Surface Enhanced Raman Spectroscopy (SERS) and attributed to the oxidation of adsorbed acetyl: *C(CH3)O *CO + *CH  *CO  CO2 [31]. It is generally believed that *C(CH3)O is an important precursor for C-C cleavage into *CO and *CHx [31-36]. However SERS also detected the electrooxidation of *CH to *CO at low potentials [31]; contrary to other opinions that *CHx are surface inhibitors and are oxidized in the second peak (~1.05V) [50, 57]. In order to identify the species contributing to the second CO2 peak, we carefully analyzed the published data on desorbed products and surface intermediates, together with our own experimental observations and the rudimentary understanding from the previous chapter. The optimal reaction pathway for the complete oxidation of ethanol to CO2 is then proposed based on the new insight into the reactions behind these two peaks. 63 Chapter 3 Fig. 3.1. DEMS mass intensities of 12CO2 and 13CO2 during oxidative stripping of adsorbed residues from isotopically labelled ethanol (a [51]), acetaldehyde (b [35]) 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 Chapter 3 3.2 Proposed Mechanisms for the Complete Oxidation of Ethanol and Acetaldehyde Peak potential is an indication of the relative difficulty in the oxidation of various adsorbed species. We propose that the second CO2 peak at higher potentials may be attributed to the oxidation of carbon-containing residues with oxygen atom(s) bound to the surface (*O-carbon species for short, e.g. adsorbed carboxylate *OCRO*). This reaction is clearly more difficult than the relatively easy *CO oxidation in the first CO2 peak. The increased difficulty in oxidizing the *O-carbon species can be due to i) a stronger adsorption of these species than *CO at higher potentials; and ii) a more difficult O-addition reaction when the C atom(s) in the *O-carbon species is not bound to Pt. The spread of the second CO2 peak over a wider potential range also indicates variety in the adsorbed species with different oxidizabilities. For example, *OC(CH3)O* would be more difficult to oxidize than *OCHO* and *O*OCCO*O* because the C–C bond in *OC(CH3)O* is orthogonal to the surface and keeps the CH3 group furthest away. The proposed complete oxidation scheme is illustrated in Scheme 3.1. The supporting arguments and evidence for it are presented summarily in §3.3. The proposed mechanism suggests that the difficulty to completely oxidize ethanol or acetaldehyde depends on the sequence of C-C cleavage and the first O-addition reaction. Complete oxidation is easier if C-C bond is cleaved before the first O-addition reaction since the resulting *CO and *CHx are relatively easy to oxidize. If C-C bond is cleaved after the first O-addition reaction, complete oxidation will involve the *O-carbon species which are difficult to oxidize. 65 Chapter 3 Scheme 3.1 The proposed pathways (non-elementary steps) for the complete oxidation of C2H5OH and CH3CHO to CO2 in different potential regions. The first O-addition to β-C involving *OCHO* or/and *O*OCCO*O as intermediates is less detrimental to subsequent complete oxidation, compared with the first O-addition to α-C involving *OC(CH3)O*, since the *OC(CH3)O* species is more difficult to oxidize as illustrated before. Nevertheless, all the complete oxidation pathways for ethanol are 66 Chapter 3 much more difficult than the direct O-addition pathways forming CH3COOH and CH3CH(OH)2 (Chapter 2). Hence, the complete oxidation pathways do not have significant contributions to current generation in ethanol solution and exhibit very low selectivity in the overall oxidation scheme. 3.3 Supporting Evidence for the Proposed Origin of the second CO2 Peak  During the oxidative stripping of residues from ethanol and acetaldehyde adsorption, *CO oxidation is generally delayed till ~ 1.0V (Fig. 3S.1 in Supporting Information 3S1). This is an indication of the difficulty in *CO oxidation.  The stripping experiment (Fig. 3S.1) also detected the presence of *OC(CH3)O* at 1.16V. *OC(CH3)O* can compete with *OH formation and could be the cause for the protracted *CO oxidation (Supporting Information 3S1). Nevertheless, since CO2 is the only desorbed product in oxidative stripping as detected by DEMS, the persistence of *OC(CH3)O* at such high potentials suggests *OC(CH3)O* as one of intermediates contributing to the 2nd CO2 peak. Besides, our own experiment also showed slow CH3COOH electrooxidation at potentials > 1.0V. (Fig. 3S.2 Supporting Information 3S2)  There are a few observations which can be used as indirect evidence for the involvement of *OCHO* and *O*OCCO*O* in the CO2 peak at ~1.0V: i) likely presence of *OCHO* and *O*OCCO*O* during stripping; ii) possibility of a CH3O* intermediate contributing to *OCHO* formation; iii) possibility of a CH3CH2O* intermediate contributing to *O*OCCO*O* formation; iv) correspondence between 67 Chapter 3 the peak potential of oxalic acid oxidation and the second CO2 peak; v) correspondence between the onset potentials of *OCHO* and *O*OCCO*O* oxidation and the 2nd CO2 peak. The discussion of these supporting observations and the development of the proposed mechanism are given in Supporting Information 3S3. 3.4 Conclusion The complete oxidation pathways for ethanol and acetaldehyde are summarized in Scheme 3.1. The selectivity of these pathways depends on whether C-C bond cleavage occurs before or after the 1st O-addition to the adsorbed residue; which forms the first and second CO2 peaks respectively. If the C-C bond cleavage is after the 1st O-addition, the difficulty and selectivity of complete oxidation also depend on whether the 1st O-addition is to α-C or β-C. Table 1 below summarizes how the electrooxidation pathways of ethanol and acetaldehyde are affected by Pt&O, Pt&α-C and Pt&β-C interactions as an extension of the understanding first established in Chapter 2. 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. Interaction Strength Strong Pt&α-C Moderate or Weak Pt&α-C Pt&O> Pt&β-C 1st *OH + α-C *C(CH3)O *OC(CH3)O*  :C(CH3)OH  CH3COOH *OC(OH)O* + *OCH3 2CO2 *CH(CH3)OH  CH3CH(OH)2 Pt&O> Pt&β-C 1st *OH + β-C *C(CH3)O *CO + *OCH3 *CO + *OCHO*  CO2 CH3CH2O* *OCH2CH2O* *O*OCCO*O* CO2 Pt&β-C >Pt&O *C(CH3)O *CO + *CH  2*CO  CO2 (Pt-Rh-SnO2) [65]: CH3CH2O* *CH2CH2O*  CO2 68 Chapter 3 A weak Pt&β-C interaction is the major cause for incomplete electrooxidation. A good catalyst for complete electrooxidation at lower potentials should therefore promote β-C adsorption either as *C(CH3)O *CO + *CH (enhancing both Pt&α-C and Pt&β-C interactions) or as C2H5O*  *CH2CH2O* (enhancing Pt&β-C interaction and weakening Pt&α-C interaction), and has the ability to effectively electrooxidize the residues formed as such. 69 Chapter 3 CHAPTER 3 – SUPPORTING INFORMATION 3S Other than the more commonly known (and discussed) intermediates of *CO, *C(CH3)O, and *CHx, adsorbed species such as *OC(CH3)O* [68], CH3CH2O* [51], :C(CH3)OH [51] have also been found. Besides, *CH(CH3)O* or CH3CHO*, i.e. η2 or η1 acetaldehydes which bond through the π orbitals, could also be likely adsorbed species [33, 68]. Similar adsorption structures have also been reported for C3 and C4 alcohols [69-72]. 3S1. Protracted *CO Electrooxidation in the Presence of Adsorbed Acetate Some important observations that can be used for deducing the origin of the 2nd CO2 peak came from the C2H5OH and CH3CHO electrooxidation studies of Shao and Adzic (Fig. 3S.1 for C2H5OH) [68]. Fig. 3S.1a shows that a small fraction of *CO from C2H5OH adsorption was more difficult to remove; requiring higher potentials than the oxidative stripping of *CO from CO adsorption. Their removal could contribute to the activity in the inception region of the second CO2 peak (< 1.0V). The long *CO tail at high potentials could be caused by the continual formation of *CO from pre-adsorbed C2 species such as *C(CH3)O and *CH(CH3)O*, and/or the delay in *OH formation (required for CO2 formation) due to the presence of some tenaciously held surface species resistant to oxidation. 70 Chapter 3 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. [68] Surface enhanced infrared absorption spectroscopy (SEIRAS) was used to identify the surface species in the oxidative stripping of C2H5OH residues (Fig. 3S.1b). The formation of *OC(CH3)O* (as identified by OCO stretching at ~1400 cm-1; and -CH3 bending at ~1350 cm-1 [32, 48, 73]) was implicated at an onset potential around 0.86V. The 71 Chapter 3 *OC(CH3)O* intensity (~1400 cm-1) in the clean electrolyte increased with potential and was maximum at around 1.06V. Unlike the case of the electrooxidation of bulk C2H5OH, there was no C=O stretching peak attributable to dissolved CH3COOH or CH3CHO (~1705-1725 cm-1). Hence the *OC(CH3)O* intermediate at high potentials was strongly surface bound and difficult to desorb. Similarly in the stripping of CH3CHO residues, the presence of *OC(CH3)O* was detected from 0.66V onwards. In fact, *OC(CH3)O* has been proposed as a catalyst poison [32, 49] which suppresses *OH formation (Regions 3 in Fig. 3S.2). Hence it has the capability to interfere *CO oxidation. 3S2. Evidence for *OC(CH3)O* Electrooxidation The electrooxidation of *OC(CH3)O* on Pt is generally believed to be infeasible at room temperature [32, 34, 48, 49]. However, since CO2 is the only desorbed product detected in the oxidative stripping of a wide range of compounds from alcohols to aldehydes, it is reasonable to propose that the *OC(CH3)O* in Fig. 3S.1b can be slowly oxidized at higher potentials. Fig. 3S.2 shows the steady state (>50 scans) cyclic voltammograms of a commercial ETEK Pt/C catalyst in 0.1M HClO4 with different CH3COOH concentrations and anodic limits of 1.17V (a) and 1.47V (b). The forward scan can be divided into five regions. The attributions for regions 1 to 4 are in accordance with earlier publications [32, 48]. For region 5 at potentials > 1.0V, the oxidation current was higher in the presence of CH3COOH than without it. This is the evidence for *OC(CH3)O* electrooxidation. 72 Chapter 3 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 Chapter 3 The following is the description of these regions. (1) The desorption of H* and adsorption of CH3COOH as *OC(CH3)O* donate electrons to the electrode with the release of protons. The total oxidation current is higher in CH3COOH due the added contribution from the dehydrogenative adsorption of the latter. (2) Since CH3COOH adsorption as *OC(CH3)O* occurs via Pt&O interaction which is enhanced by the increase in potential [48] , the electrooxidation current is slightly higher than what is normally contributed by the double layer. (3) The adsorbed *OC(CH3)O* inhibits H2O* coverage and delays the onset of *OH formation. The extent of inhibition increases with increasing CH3COOH concentration. The *OC(CH3)O* coverage may begin to decrease after the onset of *OH formation (~0.8V in 10mM CH3COOH, corresponding to the potential where *OC(CH3)O* coverage is the highest [48]). (4) Surface O* formation is also delayed, although the effect is more moderate than in region 3. (5) The *OC(CH3)O* coverage decreases further with potential up to the 1.5V upper anodic limit [74]. The increase in current density in the forward scan relative to the HClO4 baseline (Fig. 3S.2b) suggests *OC(CH3)O* oxidation as a contributing mechanism. Since O* and *OH formation on a surface partially covered with *OC(CH3)O* cannot be more facile than in the absence of CH3COOH, and the decrease in the *OC(CH3)O* coverage by desorption requires electron transfer from the electrode 74 Chapter 3 which is opposite to the observation of an oxidative current, the current increase at 1.0V and beyond must therefore involve the oxidation of *OC(CH3)O*. However, the onset of *OC(CH3)O* oxidation may begin as early as 0.8V together with *OH formation. 3S3. The Central Region of the second CO2 Peak via *OCHO* and *O*OCCO*O* The association of the central region of the second CO2 peak with *OCHO* and *O*OCCO*O* in Scheme 3.1 is supported by the following observations. (1) Likely presence of *O*OCCO*O* and *OCHO* at high potentials *O*OCCO*O* and *OCHO* absorb at around 1350 cm-1 [75] and 1320-1330 cm-1[2] respectively. As such they too may contribute to the absorption band at ~1350 cm-1 in Fig. 3S.1b which was assigned by the authors to –CH3 bending in *OC(CH3)O*. Furthermore, the intensity ratio of the 1400 cm-1 band (assigned to OCO stretching of *OC(CH3)O*) to the 1350 cm-1 band (assigned to -CH3 bending) in Fig. 3S.1b was potential dependent, suggesting that these two bands were of different adsorbed species. Since the intensity of the ~1400 cm-1 band was maximum at about 1.06V, the more significant presence of the 1350 cm-1 band at potentials slightly lower (0.96V) or higher (1.16V) is indication of other adsorbed species which absorb in the same wavelength region, i.e. *O*OCCO*O* and *OCHO*. 75 Chapter 3 (2) Possible CO2 formation from CH3O* From stripping studies using C13-labelled 12CH313CH2OH and 12CH313CHO (Fig. 3.1 with m/z = 44 indicating [12CO2]+and m/z =45 indicating [13CO2]+), the electrooxidation of the β-C12 to 12CO2 clearly contributed to both the sharp first CO2 peak and the much broader second CO2 peak, indicating that there are at least two pathways for the oxidation of β-C. The first peak is likely due to *12CO oxidation (*12CO comes from oxidizing *12CHx from the C-C cleavage of *13C(12CH3)O) [31]. The second peak has to involve another pathway and other adsorbed species with β-C12. The number of electrons released per molecule of CO2 formed (n) during CH3CHO stripping provides another valuable source of information. The isotope study of Fig. 3.1 indicates that with the increase in Ead, n increases and the CO2 from β-C gradually becomes more than the CO2 from α-C, e.g. n was 3.7 at Ead = 0.6V where the CO2 from α-C was about half of that from β-C (i.e. α-C/β-C = 0.5) [35, 36]. The greater contribution from β-C than α-C with high Ead (e.g. 0.6V) confirms that there must be some β-C1 adsorbed species more persistent than the *CO from α-C. Table 3S.1 shows the calculations implicating the likely presence of β-C1 adsorbed species. For example, CH3O* (n=5) in Scheme 3.1 (Chapter 3) could be one of the β-C1 intermediates more strongly held than *CO (n=2) from α-C. In short, if no C2 adsorbate was involved the ratio of α-C/β-C = 0.5 should give rise to an overall n of 5 x (1/1.5) + 2 x (0.5/1.5) = 4, close to the reported n = 3.7. With the presence of *OC(CH3)O* (n=3.5) and perhaps some *OCH2CH2O* (n=4), an overall n of 3.7 is possible. 76 Chapter 3 Table 3S.1 Deduction of possible β-C1 adsorbed species for the values of n = 3.7, α-C/βC = 0.5 measured at Ead = 0.6V Assuming no CO2 from adsorbed C2 species and no *CO from β-C Reactions n α-C / total C β-C / total C *CO + H2O  CO2 + 2e- + 2H+ 2 0.5/(1+0.5) 0 *OCH3 + H2O  CO2 + 5e- + 5H+ 5 0 1/(1+0.5) Overall n = 2 x 0.5/(1+0.5) + 5 x 1/(1+0.5) = 3.5 ~ 3.7 With some CO2 from adsorbed C2 species *OC(CH3)O* + 2 H2O  2CO2 + 7e- + 7H+ - + *OCH2CH2O* + 2 H2O  2CO2 + 8e + 8H n = 7/2 = 3.5 n = 8/2 = 4 Overall n = 3.7 is reasonable (3) Possible CO2 formation via CH3CH2O* Similar to C2H5OH and CH3CHO, the oxidative stripping of higher alcohols with larger alkyl fragments (Fig. 3S.3) also shows two CO2 peaks with maximums at around 0.65V 0.85V and 1.0V - 1.1V respectively. In addition the second peak to first peak ratio increases with the number of C atoms of the primary alcohol, i.e. ethanol < 1-propanol < 1-butanol (Fig 3.1a, Fig 3S.3a,c). This observation suggests that the first CO2 peak is more related to *CO from α-C while the second CO2 peak is derived mainly from the alkyl fragment. The first CO2 peak is also lower in secondary or tertiary alcohols (Fig. 3S.3 b, e, f) relative to primary alcohols with same number of carbon atoms. This can be related to the difficulty in *CRO adsorption from secondary or tertiary alcohols, resulting in low selectivity to *CO formation (the contributor to first CO2 peak) from *CRO. 77 Chapter 3 Fig 3S.3. DEMS mass intensities of CO2 formation from the oxidative stripping of 1propanol (a [69]), iso-propanol (b [70]), and four butanol isomers (c-f [72]) pre-adsorbed at various potentials. 78 Chapter 3 For tertiary butanol which adsorbs mainly as RO*, the second peak is only slightly lower than other butanol isomers. This implies that the O-addition to non-α-C in RO* and subsequent oxidation to CO2 at the potential of the second peak is possible. Extending the same reasoning to CO2 formation from ethanol, CH3CH2O* is therefore a possible intermediate. (4) Correspondence with the peak potential of oxalic acid electrooxidation 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 Chapter 3 Contrary to the electrooxidation of *CO (1.0V, 3S2), the correspondence between the peak potential of the electrooxidation of dilute oxalic acid (Fig. 3S.4) and the 2nd peak potential in oxidative stripping (Fig. 3.1, Fig. 3S.3) is supportive of (or at least not in conflict with) the involvement of oxalate oxidation during the oxidative stripping of adsorbed residues from ethanol or acetaldehyde. (5) Correspondence with the onset potential of *OCHO* and *O*OCCO*O* oxidation Unlike CH3COOH, the onset of the electrooxidation of adsorbed oxalate (*O*OCCO*O*) is earlier, which was reported to be around 0.7V [75, 76]. Although the voltammogram of the electrooxidation of 1mM oxalic acid (Fig 3S.4) also shows some delay in *OH formation, its electrooxidation is much easier than CH3COOH. The same also applies to HCOOH oxidation. Although *OCHO* is also strongly adsorbed and known to compete with *OH formation at high potentials [77], there is still some current ([...]... 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... “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... 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. .. 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. .. 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 2 Chapter 1 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. .. providing ondemand 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... 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... 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... 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 difference between the applied potential (V) and the equilibrium potential (E) of a half -cell reaction η=V–E 7 Chapter 1 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