Methanol Electrochemical Conversion to Formaldehyde over Bulk ...

Methanol Electrochemical Conversion to Formaldehyde over Bulk Metal and 

Supported Catalysts 

by 

Mohsina Islam, B.Sc., M.Sc., M.S. 

A DISSERTATION 

IN 

CHEMISTRY 

Submitted to the Graduate Faculty 

of Texas Tech University in 

Partial Fulfillment of 

the Requirements for 

the Degree of 

DOCTOR OF PHILOSOPHY 

Approved 

Carol Korzeniewski 

Chairperson of the Committee 

Shaorong Liu 

Dominick J. Casadonte Jr. 

Accepted 

Dean of the Graduate School 

May, 2006

Copyright 2006, Mohsina Islam 

ii

ACKNOWLEDGMENTS 

I would like to express my deep appreciation to my research advisor Dr. 

Carol Korzeniewski, Professor of Chemistry. It is due to her scholastic guidance, 

strong support and encouragement that ultimately made this work possible. I 

always admire and respect her as a research mentor. Under her supervision, I 

have learned many scientific and precious aspects of research and this will be a 

good guideline for my future career. I also would like to thank Dr. Dominick J. 

Casadonte Jr. and Dr. Shaorong Liu for their assistance and valuable comments 

throughout my graduate studies. 

I would like to express my appreciation to all my present and past group 

members for their assistance. Special thanks to Mr. Brandon J. Sheehan and Mr. 

Duan C. Hindes for their ingenuous assistance in machining some very 

challenging work in this research. 

I like to thank my parents for their continuous support and encouragement. 

Thank U, Mom for your support and help especially for taking care of my little 

daughter during my busy time. Finally, I am especially thankful for the patience 

and support of my husband, SM Rahmat Ullah and lots of inspiration from my 

daughter Nosheen S. Ullah. 

iii

TABLE OF CONTENTS 

ACKNOWLEDGMENTS iii 

ABSTRACT vii 

LIST OF TABLES & SCHEME ix 

LIST OF FIGURES x 

CHAPTER 

I. INTRODUCTION 

Basic Fuel Cell Concepts 1 

Polymer Electrolyte Membrane Fuel Cell (PEMFC) 4 

Direct Methanol Fuel Cell (DMFC) 6 

Methanol Oxidation Pathways 12 

Bimetallic Catalysts for Methanol Oxidation 15 

Project Focus 17 

Fluorescence Assay for Formaldehyde 19 

Summary 20 

References 

II. INSTRUMENTATION AND GENERAL EXPERIMENTAL 

CONDITIONS 

Reagents and Catalysts Materials 26 

Sonochemistry for Catalysts Synthesis 26 

Bulk Electrode Preparation 28 

iv 

1 

22 

26

Electrochemistry 30 

Formaldehyde Detection 36 

References 

III. METHANOL ELECTROCHEMICAL CONVERSION TO 

FORMALDEHYDE OVER BULK METAL ELECTRODES 

Introduction 41 

Experimental 42 

Formaldehyde Determination 43 

Results and Discussion 43 

Conclusions 61 

References 70 

IV. METHANOL ELECTROCHEMICAL CONVERSION TO 

FORMALDEHYDE OVER SUPPORTED CATALYST 

MATERIALS 

Introduction 72 

Experimental 73 

Formaldehyde Determination 73 



References 

V. A MICRO-VOLUME ELECTROCHEMICAL CELL FOR THE 

STUDY OF METHANOL ELECTROOXIDATION PATHWAYS 

v 

40 

41 

72 

100 

102 

Introduction 102

Experimental and Instrumentation 103 

Reagents 103 

Small Volume Electrochemical Cell 105 

Choice of Cell Material and Cell Fabrication 105 


Cell Characterization in 0.1M H2SO4 108 

Cell Characterization using a Reversible Redox Active 

Probe 

vi 

110 


References 

VI. SUMMARY 119 

118 

References 128

ABSTRACT 

The electrochemical oxidation of 1.0 M CH3OH in 0.1 M H2SO4 over 

different types of platinum-ruthenium (PtRu) materials was investigated. Focus 

was on the determination of formaldehyde (H2CO) produced as a function of Ru 

content in arc-melted bulk alloys and nanometer-scale catalyst materials. A 

sensitive fluorescence assay for formaldehyde, which had a detection limit down 

to 100 nM H2CO, was employed. The reaction potentials, reported in volts 

versus the reversible hydrogen electrode reference (VRHE), were in the range of 

0.5 VRHE to 0.8 VRHE. The lower potentials approach voltages that have 

technological relevance to fuel cells. Electrolysis was performed on 50 µL 

samples for a period of 180 s. Based on the coulometry, the reactant depletion 

in the cell is below 1%. 

In experiments with bulk PtRu alloys, three samples with respective Ru mole 

fraction (XRu) of 0.1, 0.3 and 0.9 were employed. Reactions were also carried out 

on bulk polycrystalline Pt for reference. Compared to Pt, the H2CO yields were 

lower for the oxidation of methanol over PtRu. Among the PtRu alloys, the H2CO 

yields decreased with increasing XRu, except at the lowest potential studied (0.5 

VRHE), where the formaldehyde yield was lowest for the sample with XRu = 0.3. 

The finding is consistent with XRu = 0.3 being the most active PtRu composition 

for methanol electrochemical oxidation at ambient temperature. The results are 

attributed to lower reactivity of methanol on the electrode with XRu = 0.9 at 0.5 

vii

VRHE due to due to blocking of initial dissociative chemisorption steps by inhibiting 

oxides present at Ru sites. 

Compared to bulk electrodes, methanol oxidation over nanometer-scale 

catalyst resulted in H2CO yields below 10 % under the conditions studied. The 

following catalyst materials were used on gold electrode: Pt-Black, C/Pt, 10 wt % 

on Vulcan XC-72R carbon, PtRu black with XRu = 0.5, PtRu catalyst prepared via 

a sonochemical (SC) method with XRu = 0.1, 0.25 and 0.5. High current densities 

were achieved during sample electrolysis. The results indicate the nanometer- 

scale catalyst converts methanol to more complete oxidation products. The 

findings are consistent with earlier studies and are attributed to readsorption and 

complete oxidation of H2CO within multiple catalyst layers, leading to lower H2CO 

yields. Similar to results for smooth, bulk electrodes, the H2CO yield was 

significantly higher for methanol oxidation over the lowest Ru content nanometer- 

scale catalyst (XRu= 0.1) and approached the response for Pt black. 

This project also advanced the design of the small volume electrolysis cell 

employed for the thesis research by incorporating a machinable MACOR glass 

ceramic disk window, which resists oxygen permeation. The cell response was 

characterized by observing the reversible, diffusion controlled waves for 

hexamine ruthenium trichloride (Ru (NH3)6Cl3) in cyclic voltammetry 

measurements. 

viii

LIST OF TABLES 

1.1 Fuel cells and their characteristics and applications. 

2.1 Sample composition in atomic % based on X-ray 

fluorescence and X-ray Diffraction. 

3.1 Summary of charge and formaldehyde yields from the 

oxidation of 1.0 M CH3OH on bulk Pt and PtRu electrode 

materials. 


oxidation of 1.0 M CH3OH over Pt based nanoscale catalyst 

materials. 


oxidation of 1.0 M CH3OH over PtRu based nanoscale 

catalyst materials. 

LIST OF SCHEME 

I Mechanism suggested for methanol oxidation to H2CO. 

ix 

5 

31 

53 

81 

85 

14

LIST OF FIGURES 

1.1: Conventional H2/O2 PEM fuel cell. 

1.2 (a): Structure of Nafion. 

1.2 (b): Three-phase model of Nafion; a fluorocarbon region (A), an 

interfacial zone (B) and an ionic cluster region (C). 

1.3: Schematic diagram of a DMFC. 

1.4: Some pathways of methanol electrochemical oxidation. 

1.5: Scheme of methanol oxidation on Pt showing the consecutive 

stripping of hydrogen atoms. 

1.6: Methanol oxidation reactions on Pt. 

1.7: Methanol oxidation reactions on PtRu. 

2.1: Scanned photographs of bulk PtRu disks. (a) XRu= 0.1, (b) XRu= 

0.3. 

2.2: Front portion of micro-volume electrochemical cell used for 

sample electrolysis. 

2.3: Block diagram of the automated formaldehyde analyzer. 

2.4: Calibration curve for standard H2CO solution. Concentration in 

the range of 2x10 -6 -2.5x10 -5 M. 

2.5: Fluorescence signal of standard H2CO solution. 

3.1: Cyclic voltammograms of bulk electrode materials employed in 

the study. Scans were recorded in Ar purged 0.1 M H2SO4. 

Polycrystalline Pt electrode (Pt-solid, a); bulk PtRu alloys (PtRusolid) 

with Ru mole fraction (XRu) of 0.1 (b), 0.3 (c) and 0.9 (d). 

Cyclic voltammetric measurements were performed by scanning 

the potential between 0.05 to 0.75 VRHE. The scan rate was 50 

mV/s. 

3.2: Cyclic voltammogram of a polycrystalline Pt electrode in Ar 

purged 0.1 M H2SO4 recorded at a scan rate of 50 mV/s. 

3.3: H2CO yields from methanol electrochemical oxidation on a 

polycrystalline Pt electrode. All experiments involved the 

x 

7 

8 

8 

11 

13 

18 

18 

18 

32 

34 

38 

39 

39 

47 

48 

54

electrolysis of 1.0 M CH3OH in 0.1 M H2SO4 in a volume of 50 µL 

for a period of 180s. 

3.4: H2CO yields from methanol electrochemical oxidation on a PtRu- 

solid (XRu = 0.1) electrode. All experiments involved the 











3.7: Comparison of H2CO yields from methanol (1.0 M) 

electrochemical oxidation on bulk electrode materials at 0.6 and 

0.5 VRHE. 

3.8: Plot of the nanomoles H2CO formed vs. the reaction charge 

passed following the oxidation of 1.0 M methanol on a 

polycrystalline Pt electrode in 0.1 M H2SO4 (Reaction time = 180 

s). 

3.9: Plot of H2CO formation rate averaged over 180 s electrolysis 

periods vs. potential for the oxidation of 1.0 M methanol in 0.1 M 

H2SO4 on a polycrystalline Pt electrode. 


passed following the oxidation of 1.0 M methanol on a bulk alloy 

PtRu- solid (XRu = 0.1) electrode in 0.1 M H2SO4 (Reaction time = 

180 s). 



H2SO4 on a bulk alloy PtRu- solid (XRu = 0.1) electrode. 




180 s). 


xi 

67 

55 

56 

57 

58 

62 

63 

64 

65 

66






180 s). 




4.1: Cyclic voltammogram of a polycrystalline gold electrode in 0.1 M 

H2SO4; Scan rate: 50 mV/s. 

4.2: Cyclic voltammogram of a polycrystalline gold electrode in 0.1 M 

H2SO4 after adsorption of C/Pt, 10 wt% (a) or Pt-black (b) catalyst 

films; Scan Rate: 50 mV/s. Arrow represents the stripping wave 

for gold oxide. 

4.3: Cyclic voltammograms of catalyst materials employed in the 

study. Scans were recorded in Ar purged 0.1 M H2SO4 at 50 

mV/s. The following nanoscale catalyst materials were studied as 

thin films adsorbed on a Au electrode- Johnson Matthey (JM) 

PtRu with XRu = 0.5 (a); PtRu catalyst prepared via a 

sonochemical (SC) method with XRu = 0.1 (b) and 0.5 (c). 

4.4: H2CO yields from methanol electrochemical oxidation in 0.1 M 

H2SO4 on Pt -Black catalyst on a polycrystalline Au electrode. All 

experiments involved the electrolysis of 1.0 M CH3OH in 0.1 M 

H2SO4 in a volume of 50 µL for a period of 180s. 


H2SO4 on C/Pt, 10 wt% catalyst on a polycrystalline Au electrode. 

All experiments involved the electrolysis of 1.0 M CH3OH in 0.1 M 

H2SO4 in a volume of 50 µL for a period of 180s. 


H2SO4 on SC PtRu, XRu= 0.1 catalyst on a polycrystalline Au 

electrode. All experiments involved the electrolysis of 1.0 M 

CH3OH in 0.1 M H2SO4 in a volume of 50 µL for a period of 180s. 


H2SO4 on SC PtRu, XRu= 0.1 catalyst, capped by Nafion on a 

polycrystalline Au electrode. All experiments involved the 


xii 

68 

69 

76 

77 

79 

82 

82 

86 

86











H2SO4 on JM PtRu, XRu= 0.5 catalyst on a polycrystalline Au 



4.11: H2CO yield from methanol (1.0 M) electrochemical oxidation on 

catalyst materials at 0.5 VRHE. 



H2SO4 over a Pt -Black catalyst layer adsorbed on a 

polycrystalline Au electrode. 



H2SO4 over a C/Pt, 10 wt% catalyst layer adsorbed on a 




H2SO4 over a SC PtRu, XRu= 0.1 catalyst layer adsorbed on a 




H2SO4 over a SC PtRu, XRu= 0.1 catalyst layer (capped by 

Nafion) adsorbed on a polycrystalline Au electrode. 






xiii 

95 

87 

87 

88 

89 

93 

93 

94 

94 

95






H2SO4 over a JM PtRu, XRu= 0.50 catalyst layer adsorbed on a 



passed following the oxidation of 1.0 M methanol in 0.1 M H2SO4 

over a Pt-Black catalyst layer adsorbed on a polycrystalline Au 

electrode (Reaction time = 180 s). 



over a C/Pt, 10 wt% catalyst layer adsorbed on a polycrystalline 

Au electrode (Reaction time = 180 s). 



over a SC PtRu, XRu= 0.10 catalyst layer adsorbed on a 

polycrystalline Au electrode (Reaction time = 180 s). 



over a SC PtRu, XRu= 0.10 catalyst layer (capped by Nafion) 

adsorbed on a polycrystalline Au electrode (Reaction time = 180 

s). 











over a JM PtRu, XRu= 0.10 catalyst layer adsorbed on a 


xiv 

96 

96 

97 

97 

98 

98 

99 

99

5.1: 

Exploded view of electron transfer at an electrode surface in a 

thin-layer flow cell. 

5.2: Side view and top view of the electrochemical cell window. 

5.3: Cyclic voltammograms of a polycrystalline Pt electrode in 0.1 M 

H2SO4 recorded in a conventional electrochemical cell (a) and the 

electrolysis cell with the electrode mounted far from MACOR 

window (b). The scan rate was 50 mV/s. 


H2SO4 recorded in an electrolysis cell with the electrode mounted 

near the MACOR window. Scan rate was 10 mV/s (a) and 50 

mV/s (b). 


H2SO4 recorded in an electrolysis cell with the electrode mounted 

near the MACOR window (a) and mounted inside the MACOR 

window (Arrow represents skewing of the hydrogen adsorption 

region) (b). The scan rate was 10 mV/s. 

5.6: Cyclic voltammograms of [Ru(NH3)6Cl3] in 0.1 M NaNO3 recorded 

in a conventional electrochemical cell (a, b) and the MACOR 

cavity of the thin layer electrolysis cell shown in Figure 3.2 (c, d). 

The [Ru(NH3)6] 3+ concentration and scan rates were : (a) 5 × 10 -3 

M and 50 mV/s , (b) 5 × 10 -3 M and 0.2 mV/s (c) 2.5 × 10 -3 M and 

0.2 mV/s (d) 2.5 ×1 0 -2 M and 1 mV/s. The scans started from 0.4 

VAg/AgCl and swept toward more negative potentials. 

6.1: Bar charts comparing the formaldehyde yields from reaction of 

1.0 M CH3OH in 0.1 M H2SO4 during 180 s electrolysis periods for 

solid bulk electrodes. 

6.2: Bar charts comparing the formaldehyde yields from reaction of 

1.0 M CH3OH in 0.1 M H2SO4 over nanoscale catalyst materials 

during 180 s reaction periods. 

xv 

104 

107 

111 

112 

113 

116 

123 

124

CHAPTER I 

INTRODUCTION 

Basic Fuel Cell Concepts 

Fuel cells have emerged as a promising technology for electric power 

generation. Fuel cell development spans more than 150 years. This long period 

has been characterized by many ups and downs [1]. In 1839, William Robert 

Grove first demonstrated the fuel cell using a simple experiment where a reverse 

of water electrolysis was performed to generate electricity from hydrogen (H2) 

and oxygen (O2) gases. Grove stated a definition for this type of fuel cell that still 

applies today –“A Fuel cell is an electrochemical device that continuously 

converts chemical energy into electrical energy (and some heat and water) 

without any combustion for as long as an oxidant and reactant (fuel) are 

supplied” [2]. During the 1960s, the U.S. National Aeronautics and Space 

Administration (NASA) used precursors to today’s fuel cell technology as power 

sources in Gemini and Apollo spacecraft, and fuel cells still provide electricity and 

water for the space shuttle [2]. Fuel cells are similar in many ways to batteries. 

They both share the electrochemical nature of the power generation process. 

Furthermore, like internal combustion engines (ICEs), fuel cells will operate 

continuously by consuming some type of fuel [2]. 

The fuel cell consists of an anode and a cathode separated by an 

electrolyte membrane. In a hydrogen-oxygen fuel cell (Figure 1.1), hydrogen 

flows past the anode, which contains a catalyst coating that facilitates oxidation 

1

of the gas to form protons (H + ) and electrons (e - ). The electrolyte membrane 

passes the protons to the cathode side of the fuel cell, and the electrons released 

flow through the external circuit. At the cathode, a catalyst coating facilitates the 

reduction of oxygen: protons and electrons combine with oxygen atoms on the 

catalyst surface to produce water (Equations 1-3) [2]. 

Anode reaction: H2 (g) → 2H + (aq) +2e - (1) 

Cathode reaction: ½ O2 (g) + 2H + (aq) + 2e - → H2O (l) (2) 

Cell reaction: H2 (g) + ½ O2 (g) → H2O (l) (3) 

The oxygen required for fuel cell operation typically comes from air, but 

hydrogen is not as readily available. Hydrogen can be produced from different 

processes, including natural gas steam reforming, electrolysis, photoelectrolysis, 

biomass gasification and pyrolysis and photobiological methods. In natural gas 

steam reforming, the first step is to expose natural gas to high temperature 

steam to produce hydrogen, carbon monoxide, and carbon dioxide. The second 

step is to convert the carbon monoxide with steam to additional hydrogen and 

carbon dioxide. In electrolysis, electric energy is used to split water to form 

hydrogen and oxygen gas. Photoelectrolysis is a process where sunlight 

absorbed by a semiconductor electrode provides the energy to split water into 

hydrogen and oxygen. The production of hydrogen can result from high 

temperature gasification and low temperature pyrolysis of biomass (i.e., wood 

chips and agricultural wastes). Photobiological processes produce hydrogen 

using light energy through the metabolic activities of certain photosynthetic 

microbes. 

2

There are some limitations in the use of hydrogen. As hydrogen is a gas, 

it is difficult to store and distribute. Methanol, natural gas, coal gas, 

hydrocarbons (methane, propane) are some additional potential fuels. A 

reformer can be used for turning hydrocarbon or alcohol fuels into hydrogen, 

which is then fed into the fuel cell [3]. Alternatively, simple alcohols can be 

directly used although the reaction kinetics are slower than for hydrogen. 

Catalysts are required to enable fuel cell reactions. The cost, availability and 

performance of catalyst materials are important criteria in the development of a 

fuel cell. 

Since in a fuel cell, the energy stored in the bonds of chemical reactants is 

converted directly to electricity, in theory a fuel cell can operate at much higher 

efficiencies than ICEs, which require the intermediate production of heat from 

chemical reactants [2-4]. Fuel cells are conventionally said to be non Carnot- 

limited [1] and have potential to extract more electricity from the same amount of 

fuel compared to ICEs. Fuel cells can have zero or ultra low emissions, which 

makes them attractive for transport applications. The fuel cell itself has no 

moving parts, making it a quiet source of power and not readily susceptible to 

mechanical breakdown. The potential of fuel cells to operate efficiently lowers 

fuel consumption, which benefits conservation efforts and results in reduced 

emission of gaseous pollutants and the green house gas CO2 (the latter is 

emitted when hydrocarbons are reformed to produce H2). 

There are several types of fuel cells, which are mainly classified according 

to their electrolytes and operating temperatures. These include the polymer 

3

electrolyte membrane fuel cell (PEMFC), phosphoric acid fuel cell (PAFC), direct 

methanol fuel cell (DMFC), alkaline fuel cell (AFC), molten carbonate fuel cell 

(MCFC), and solid oxide fuel cell (SOFC). The characteristics of all fuel cells are 

summarized in Table 1.1 [1-3]. Low operating temperatures, rapid start up, light 

weight, high power density and simplicity make PEMFCs and DMFCs attractive 

for transportation and consumer electronics. 

Polymer Electrolyte Membrane Fuel Cells (PEMFCs) 

The electrolyte in PEMFCs or SPEFCs (Solid Polymer Electrolyte Fuel 

Cells) is a polymeric ion exchange material that enables high proton conductivity. 

Nafion, which is a perfluorosulfonated polymer consisting of 

polytetrafluorethylene backbone and perfluoroether side chain with a sulfonate 

group is commonly used. The PEM is a thin sheet (50-200 µm) that when 

hydrated allows hydrogen ions to cross it. A general schematic of a PEMFC is 

shown in Figure 1.1. These fuel cells are capable of operating at ambient 

pressure and temperatures just below 100 ºC. PEMFCs were the first fuel cells 

to be used in space exploration missions. The Gemini program employed a 1kW 

fuel cell stack as an auxiliary power source [5]. The historical development of 

PEM fuel cells has been described recently [6]. PEMFCs also provided the 

astronauts with clean drinking water. Initially a polystyrene sulfonate (PSS) 

polymer was employed as a membrane, which proved unstable. This led NASA 

to rely on AFC systems. 

A major breakthrough in the field of PEM fuel cells came with the 

development of Nafion membranes by DuPont. These membranes possess a 

4

Table 1.1. Fuel cells and their characteristics and applications [1-3]. 

Fuel cell 

type 

Electrolyte Charge 

carrier 

Op. 

temp. 

(˚C) 

5 

Efficien 

cy 

(%) 

Life time 

(hr) 

Power 

range/Applications 

AFC KOH OH - 60-120 35-55 >10,000 40,000 5-250 kW, 

automotive, CHP 

(combined heat 

and power 

generation) 

PAFC H3PO4 H + ~220 40 >40,000 200 kW, portable 

CHP 

MCFC Lithium CO3 

and 

potassium 

carbonate 

2- ~650 > 50 >40,000 200 kW-MW range, 

CHP & standalone. 

SOFC Solid oxide 

(ytria, 

zirconia) 

DMFC H2SO4/ 

HClO4 

O 2- ~1000 > 50 >40,000 2 kW-MW range, 

CHP & standalone. 

H + 50-100 40-50 >10,000 1-100 kW. 

Transportation, 

remote power

higher acidity, proton conductivity and far greater stability than polystyrene 

sulfonate. The structure of Nafion is shown Figure 1.2. 

As PEMFCs use solid membrane electrolyte and the only liquid product is 

water, corrosion problems are minimal. PEMFCs can produce high power 

densities and offer the advantages of low weight and small volume. The main 

disadvantage of the PEMFC is that it is highly sensitive to fuel impurities, such as 

carbon monoxide (CO), which is produced as a by-product of hydrocarbon 

reforming to hydrogen [7]. The CO level of 10-50 ppm poisons the catalyst, 

causing severe degradation of cell performance. The effect has become 

pronounced when CO covers platinum catalyst sites, preventing the hydrogen 

fuel from reacting [3, 7]. 

Direct Methanol Fuel Cells (DMFCs) 

These cells are similar to the PEMFC in that they both can use a polymer 

membrane as electrolyte. DMFCs operate at the same temperatures as 

PEMFCs, although slightly higher temperatures are preferable in order to 

improve the power density. In DMFCs, the anode catalyst directly oxidizes 

methanol, eliminating the need for a fuel reformer, which simplifies operation. 

DMFCs are being targeted for tiny to mid size applications like cellular phones 

and laptop computers. Compared to hydrogen, methanol is an attractive fuel 

option. It can be produced from natural gas or renewable biomass resources. 

Among the alcohol fuels, methanol is the most easily oxidized fully to CO2, and it 

is easy to store, transport and distribute. The DMFC can be operated with liquid 

or gaseous methanol/water mixtures. It is assumed that the existing 

6

Figure 1.1: Conventional H2/O2 PEM fuel cell 

[3]. 

7

Figure 1.2 (a): Structure of Nafion. 

Figure 1.2 (b): Three-phase model of Nafion; a fluorocarbon region (A), an 

interfacial zone (B) and an ionic cluster region (C) [5]. 

8

infrastructure for fuels may be adapted to methanol, which has the advantage of 

a high specific energy density [5]. Thus, methanol has become a potential 

candidate to power fuel cells [8, 9]. 

DMFCs were first studied over 40 years ago. At that time, an alkaline 

electrolyte was used. However, carbonation of the alkaline electrolyte (caused 

by the evolved CO2) decreased efficiency by reducing the electrolyte conductivity 

and de-polarizing the cathode [10]. In the late 1970s to early 1980s, several 

researchers explored the use of H2SO4 as an electrolyte [11]. Currently, 

researchers focus on the use of a PEM capable of transporting proton. A PEM 

has the same advantages in a DMFC as it does in a PEMFC operating on H2 and 

O2 (e.g. lightweight, good proton conductivity). 

Figure 1.3 shows a schematic diagram of a DMFC. Electrons and protons 

are liberated at the anode with the aid of a catalyst. The electrons and protons 

travel through the external circuit and PEM, respectively, to the cathode 

electrocatalyst where they are consumed together with oxygen in a reduction 

reaction (Equations 4-6) [2]. 

Anode reaction: CH3OH (aq) + H2O (l) → CO2 (g) +6H + (aq) +6e - (4) 

Cathode reaction: 6H + (aq) +6e - + 3/2 O2 (g) → 3 H2O (l) (5) 

Cell reaction: CH3OH (aq) + 3/2 O2 (g) → CO2 (g) + 2 H2O (l) (6) 

DMFCs have limitations, which include slow oxidation kinetics of methanol 

and reactant crossover through the membrane [2, 12]. Although significant 

performance gains have been achieved using PEMs, many problems still exist. 

The best available and widely used PEM to date is Nafion. However, Nafion is 

9

permeable to methanol (and water). This permeability enables methanol to 

“crossover” from the anode to the cathode and lowers cathode performance. 

Since methanol will be oxidized at the cathode, it creates a “mixed” electrode 

potential that will be lower than the standard potential for oxygen reduction. In 

order to circumvent this problem, one can modify the catalyst composition. In 

addition, several alternatives to Nafion are being investigated. One common 

approach is to create membranes that exhibit less methanol crossover. 

Membranes based on sulfonated poly (ether ketone) [13], sulfonated 

polysulfones and poly (ether sulfones) [14] have been developed and tested. 

While these membranes exhibit less methanol permeability than Nafion, their 

proton conductivities are substantially lower and hence they do not perform as 

well. Pickup et al. [15, 16] have developed composite poly (1- 

methylpyrrole)/Nafion membranes. These membranes exhibit ca. 50% less 

methanol crossover without a significant increase in the resistance (or decrease 

in conductivity) of the composite membrane. 

While methanol is the most common hydrocarbon that is being explored 

for direct oxidation in fuel cells, many other hydrocarbons have been either 

proposed or studied. Savadogo’s group has investigated running fuel cells 

directly on both propane [15] and acetals [17]. Qi, Z. has recently reported good 

low-current density performance for cells running on 2-propanol [18]. Ethanol is 

also being explored as a possible fuel [19]. 

10

Figure 1.3: Schematic diagram of a DMFC [3]. 

11

Methanol Oxidation Pathways 

The electrochemical-oxidation of methanol has attracted extensive interest 

over the past few years because of its potential in DMFC applications [20-23]. In 

a typical DMFC anode, methanol is oxidized over nanometer scale catalyst 

particles to CO2 through a multi-step six electron process (CH3OH+ H2O → CO2 

+ 6H + +6e - ). Some pathways for methanol electrochemical oxidation are shown 

in Figure 1.4. The conversion to CO2 is complicated, as several pathways are 

possible. Furthermore, the involvement of the electrode surface gives rise to 

geometric and electronic effects that are poorly understood and difficult to probe. 

The continuous electrochemical oxidation of methanol generally suffers from 

poisoning of the catalyst by adsorbed CO. CO poison formation from methanol 

and its inhibiting effect have been studied exhaustively [24, 25, 26]. It is also well 

established from early studies that methanol oxidation occurs via parallel 

pathways giving rise to a complex network of reactions involving formaldehyde 

(H2CO) and formic acid (HCOOH) as side products and sources of efficiency loss 

for DMFC operation. [9, 27, 28, 29, 30-34, 41]. It has been proposed [31] that 

the elementary steps leading to H2CO accumulation could start with the 

adsorption of a methanol molecule via the oxygen atom to form a methoxide 

species (H3CO)ad. The reaction involves the elimination of one hydrogen atom 

(as a H + ) delivering one electron to the metal [CH3OH → (H3CO)ad + H + +e - ]. In 

a second step (H3CO)ad can eliminate a second H atom from the CH3 group, 

forming H2CO which then desorbs [(H3CO)ad → H2CO + H + + e - ]. The 

mechanism discussed [31] above is shown in Scheme I. The study of these 

12

- + 

- 2e , - 2H 

- + 

CH 3OHCH 2O 

- 2e , - 2H 

HCOOH 

+ H O 

- + 

- 2e , - 2H 

- + 

- 4e , - 4H 

+ H O 

2 

CO 

- + 

- 2e , - 2H 

CO 

Figure 1.4: Some pathways of methanol electrochemical oxidation. 

13 

ads 

2 

2 

- H O 

2 

- + 

- 2e , - 2H

eactions especially pathways leading to H2CO have been unexplored until 

recently due to interferences which inhibit detection of the intermediate species 

Scheme I: Mechanism suggested for methanol oxidation to H2CO [31]. 

by modern in situ analysis techniques like mass spectrometry [35,36] and 

infrared spectroscopy [37-39]. Direct detection of H2CO by mass spectrometry is 

generally hindered by the methanol mass fragments at m/z = 28-30 [27, 35, 36]. 

IR spectroscopy is also incapable of producing H2CO spectra as H2CO in 

aqueous solution forms a gem-diol (H2C(OH)2) which overlaps with water 

vibrational bands [37-39]. Recent studies [27, 28, 30, 40] have brought serious 

attention to the stable by-products of methanol oxidation, as these affect CO2 

production and decrease energy conversion efficiency [9, 41-43]. These efforts 

mainly involved the use of differential electrochemical mass spectrometry 

(DEMS), which can only detect CO2 and methyl formate directly. In their DEMS 

studies, assuming that CO2, formic acid and H2CO are the only direct reaction 

products, the faradic current for H2CO was found indirectly as the difference 

between the total Faradic charges measured and the partial currents/charges of 

CO2 and formic acid formation, calculated using corresponding calibration 

constants [27, 28, 40]. In this research project, we are mainly interested in 

14

H2CO, which is formed in the first two electron oxidation step (CH3OH → H2CO + 

2H + +2e - ). The determination of H2CO produced from methanol oxidation over 

smooth or porous catalyst/electrode materials during short electrolysis periods 

requires very sensitive analytical methods, since small amounts of H2CO are 

produced. A recent study by the Iwasita group demonstrated H2CO 

quantification using a longer electrolysis time (1000s) and smooth polycrystalline 

and Pt(III) electrodes with high performance liquid chromatography (HPLC) as a 

detection method [31]. For the last few years, our group has emphasized a direct 

measurement technique for H2CO based on fluorescence spectroscopy [29, 41]. 

Methanol electrochemical oxidation to CO2 is complex and significant 

improvements in catalyst development are required for low-temperature DMFCs 

[9]. Conventionally, Pt based catalyst materials have been employed. The 

oxidation of methanol on Pt based catalysts is assumed to be proceeded by the 

adsorption of the molecule followed by several steps of deprotonation which 

results the main catalyst poison Pt-COads [5]. A scheme for this process is shown 

in Figure 1.5. These CO intermediates are only removed when there are 

oxygenated species present on the Pt surface. The mechanisms of these 

reactions on Pt are represented in Figure 1.6. 

Bimetallic Catalysts for Methanol Oxidation 

Since Pt surface is extremely susceptible to poisoning and rapidly blocked 

by adsorbed CO intermediates, methanol electrochemical oxidation reactions on 

Pt are comparatively slow [7, 47-49]. To enhance the catalytic activity and 

increase the progress of CO removal by oxidation to CO2, a number of co-metals 

15

have been added to Pt catalysts. Studies have shown that the most promising 

electro-catalysts are bimetallic materials consisting of Pt and a second transition 

metal including Ru, Re, Os, Mo, Pb, Bi and Sn that activates H2O and promote 

the oxidation of CO to CO2 [47-49]. Of all bimetallic materials mentioned, Pt-Ru 

has long been recognized as exhibiting superior catalytic activity, if compared to 

Pt and are by far the most widely used anode catalysts for DMFCs [9, 48, 49]. A 

bifunctional mechanism [50-52] is considered to be responsible for the 

enhancement of catalytic activity. Although Pt is an effective catalyst/electrode 

material for methanol dehydrogenation, water activation can be enhanced by 

incorporating ruthenium (Ru + H2O → Ru-OHads + H + + e - ), which efficiently 

breaks the H-OH bonds in the water molecule and forms the surface oxide 

necessary for the reaction to promote the CO2 production. It is anticipated that 

the addition of Ru lowers the potential for methanol oxidation by facilitating the 

removal of CO from the electrode surface, thereby increasing CO2 production [9, 

21, 40, 48, 49, 53-57]. The bifunctional mechanism, first proposed by Watanable 

and Motoo [50] identifies the capability of Ru to nucleate oxygen-containing 

species at ca. 300 mV more negative than Pt. Additionally, Ru is a fairly noble 

metal and therefore much more stable than other promoters under the conditions 

of DMFC operation [5]. The simplified mechanism for methanol oxidation using 

Pt-Ru is shown in Figure 1.7. 

The catalytic activity of PtRu catalyst/electrode materials is sensitive to 

Pt/Ru composition and structure. The performance of PtRu alloy catalysts with a 

certain Pt/Ru ratio in methanol oxidation reaction also depends on the electrode 

16

potential as well as on the methanol concentration and reaction temperature [40, 

51, 58]. It was observed that adding a small or medium amount of Ru (Ru mole 

fraction, XRu = ~ 0.1-0.5) to Pt electrode materials greatly increases catalytic 

activity toward methanol oxidation [32, 40, 51, 53]. Iwasita noted that a catalyst 

composition of Pt/Ru between XRu = ~ 0.1-0.5 provides adequate Pt and Ru sites 

for the efficient methanol electrochemical reaction [56]. 

Project Focus 

The goal of this project was to evaluate the importance of H2CO as a 

stable reaction by-product of methanol electrochemical oxidation, as 

accumulation of H2CO can lower efficiency in DMFCs. The study focused on the 

measurement of H2CO produced during the oxidation of methanol over Pt and 

PtRu catalyst materials, including arc-melted, bulk PtRu alloys and nano meter 

scale PtRu catalyst with varying Ru surface composition. The bulk alloys were 

prepared and characterized at the Lawerence Berkeley National Laboratory and 

have become standard reference materials for the study of CO and methanol 

electrochemical oxidation reactions in support of PEM fuel cell anode 

development. This dissertation reports the first results on H2CO production from 

methanol electrochemical oxidation over the standard materials. A sensitive 

fluorescence assay was employed that enabled H2CO to be quantified following 

short (180 s) electrolysis periods in solutions that contained 1.0 M methanol in 

0.1 M H2SO4. Reactions on bulk Pt and PtRu electrodes were studied first, and 

afterward the results were applied to understand subsequent measurements of 

17

Figure 1.5: Scheme of methanol oxidation on Pt showing the consecutive 

stripping of hydrogen atoms [5]. 

Figure 1.6: Methanol oxidation reactions on Pt. 

Figure 1.7: Methanol oxidation reactions on PtRu. 

18

methanol oxidation on alloy-like nanometer-scale catalyst particles of Pt and 

PtRu. Quantitative detection of H2CO was critical as the electrochemical 

conversion of methanol to H2CO is typically low especially with porous catalyst 

particles. Our aim was to study the H2CO yield with different catalysts and 

electrodes to identify the factors that limit catalyst and electrode performance in 

fuel cells. 

Fluorescence Assay for Formaldehyde 

The direct detection of H2CO during methanol electrochemical conversion 

is very important because H2CO can cause severe deficiency in fuel cell 

performance. Disregarding the in-situ IR spectrometry and mass spectrometry 

techniques for H2CO analysis, liquid chromatography (LC) is a possible 

technique [31]. However, the chromatographic method is slow and requires 

longer electrolysis periods to attain enough sensitivity for H2CO detection [31]. 

This method is also restricted by the requirement of bulky, expensive and 

elaborate instrumentation. For our purpose, we were interested in the simple 

and fast fluorometric flow injection analysis (FIA) method developed by Dasgupta 

and Dong [45]. The analysis utilizes the Hantzsch reaction involving the 

cyclization of an amine, an aldehyde, and a β-diketone to form a dihydropyridine 

derivative. The reaction was introduced by Nash for analytical purposes in 

photometric versions using 2, 4-pentadione [60]. Based on the Hantzsch 

reaction, the specific use of 2, 4-pentadione and ammonium acetate (which 

serves as a source of NH3 and as a buffering agent) for the determination of 

H2CO leads to 3, 5-diacetyl-1, 4-dihydrolutidine (DDL): 

19

O O 

2 + H2CO + NH3 

DDL fluoresces optimally at 510 nm and has an excitation maximum at 412 nm 

[61]. A number of reagents were studied for fluorometric determinations, and 1, 

3-cyclohexadione (CHD) and 5, 5-dimethyl-CHD (dimedone) were found reactive 

with all aldehydes [45]. It was also determined that 1, 3-cyclohexadione can 

provide substantially more sensitivity than 2, 4- pentadione with acceptable 

selectivity over other aldehydes if proper reaction time is maintained [44]. 

In the present work, 1, 3-cyclohexadione was used for automated real- 

time determination of H2CO. A flow injection analysis arrangement was 

employed (Chapter II). 

60 °C, 5 min 

Summary 

This dissertation describes investigations of H2CO production during the 

electrochemical oxidation of methanol over bulk metal electrodes and nanometer 

scale catalyst particles through the use of a micro volume electrochemical cell 

and a sensitive fluorescence assay for H2CO. In Chapter II, a description of the 

general experimental procedures and instrumentation are presented. Chapter II 

also includes a brief discussion of the electrochemical cell employed for 

methanol electrolysis. Chapters IIII and IV focus on the electrochemical 

oxidation of methanol to H2CO during reactions over Pt and PtRu solid metal 

electrodes and nanometer scale catalyst materials, respectively. The 

characterization of the electrolysis cell is described in Chapter V. Chapter VI 

20 

O O 

N 

+ 3 H2O

provides a summary which compares experimental results for bulk and catalyst 

electrode materials. 

21

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25

CHAPTER II 

INSTRUMENTATION AND GENERAL EXPERIMENTAL CONDITIONS 

Reagents and Catalysts Materials 

The following metal salts and catalyst materials were obtained from 

Johnson Matthey (JM) (Ward Hill, MA): PtRu black (50 at. % Pt, 50 at. % Ru, JM 

PtRu, XRu - 0.5); Pt at 10 wt % metal loading on Vulcan XC-72R carbon (C/Pt, 10 

wt %); Pt-black, PtBr4 (99.99 %), RuCl3. x H2O (99.9 % purity, 42 wt % Ru 

content) and hexaammine ruthenium (III) trichloride. (Ru(NH3)6Cl3, 32.6%, Ru 

content). Concentrated H2SO4 and HClO4 were obtained in 99.999% purity from 

Aldrich (Milwaukee, WI). Potassium hydroxide (Aldrich), 1, 3- cyclohexadione 

(CHD, Aldrich), ammonium acetate (Aldrich), HCl (Mallinckrodt, Hazelwood, MO) 

and sodium nitrate (Mallinckrodt chemicals) were reagent grade or better and 

used as received. Methanol (99.9 % purity, Fisher Scientific, Houston, TX) was 

washed over alumina, filtered, distilled, and stored refrigerated. Aqueous 

solutions containing methanol were prepared fresh the day of each experiment. 

Argon (Trinity Gases, Dallas, TX) was of ultrahigh purity. All solutions were 

prepared with deionized water (18 MΩ cm) from a four-cartridge Nanopure 

Infinity System (Barnstead, Dubuque, IA). 

Sonochemistry for Catalysts Synthesis 

PtRu alloy nanoparticles (SC PtRu) with various Ru surface compositions 

(XRu = 0.1, 0.25 and 0.5) were synthesized using a sonochemical (SC) method 

reported earlier [1]. The procedure is outlined here for reference. A Sonics and 

26

Materials Vibra Cell TM immersion sonicator operating with 600 W electrical input 

power and calorimetrically measured 17 W/cm 2 acoustic powers [2,3] was used. 

The optimization of the sonication cell was performed either by using KI [4] or an 

8891 ultrasound cleaning bath (Cole-Parmer Instrument Company, operating 

frequency: 47 kHz). The cell temperature was maintained at 5º C while the 

cooling bath temperature was - 20 ºC. 

In the sonochemically prepared PtRu alloy synthesis procedure, there is a 

reaction between a mixture of Pt/Ru metal salts (PtBr4 and RuCl3) and lithium 

metal occurring in the sonication cell. Prior to synthesis, the cell was dried in an 

oven at 160º C, then purged with nitrogen and charged with tetrahydrofuran 

(THF, 15 mL), lithium metal cuttings (4.32 mmol) and naphthalene (2.81 mmol). 

A solution of 15 mL THF was prepared by dissolving an appropriate amount of 

PtBr4 and RuCl3 (0.05, 0.125 mol Ru 3+ and 0.45, 0.375 mol Pt 4+ for particles with 

XRu ≈ 0.1 and XRu ≈ 0.25 respectively and 0.25 mol in each metal for particles 

with XRu ≈ 0.5) in a previously dried (at 160º C for more than 6 hrs) Schlenk flask. 

After the dissolution of both metal halides, the resulting suspension looked black 

red-brown. The cell was sonicated for 2 hrs after cooling at -20º C for 15 min. 

During the sonication of cell containing lithium metal cutting, the RuCl3/PtBr4 

mixture was slowly introduced to the dark purple-blue lithium naphthalide THF 

solution passing through a cannula. The reaction mixture bubbled during the 

addition and ultimately turned black. The sonication was continued for 4 more 

hours. 

The reaction mixture was transferred to a 250 mL beaker in air. A black 

27

precipitate was collected after centrifuging the mixture. For washing purposes, 

the precipitate was rinsed with THF until it appeared colorless and odorless. The 

precipitate was then washed with 20–25 mL of deionized water. After 5–10 min 

stirring, the powder was finally filtered and centrifuged. The centrifuge tube was 

half filled with the suspension, and then filled in full by THF. When THF was 

added to the suspension, a few bubbles were observed. A compact black color 

precipitate with yellow–green supernatant was obtained after centrifuging for 

approximately 2 min. The precipitate was repeatedly washed by THF until a 

colorless supernatant was observed. The resulting metal black was collected with 

THF (no more than 30 mL), and transferred into a vial equipped with a small 

magnetic stirring bar. The suspension was continually stirred in order to prevent it 

from nucleating. The procedure produced particles with diameters in the range of 

2–6 nm as observed by transmission electron microscopy. 

Bulk Electrode Preparation 

The following bulk materials were employed as working electrodes: 

polycrystalline platinum (Pt-solid, 0.81 cm 2 ), PtRu alloys (PtRu-solid) with Ru 

mole fractions (XRu) of 0.1 (0.71 cm 2 ), 0.3 (0.66 cm 2 ) and 0.9 (0.69 cm 2 ) and 

polycrystalline gold (Au) (0.67 cm 2 ). The preparation methods of the PtRu bulk 

alloys were described previously [5, 6]. In their work, the bulk alloys with same 

Ru compositions similar to our samples were observed to have the chemical 

composition and lattice constants which are shown in Table 2.1 [5]. The PtRu 

(PtRu-solid, XRu = 0.1 and 0.9) metal disks were received as polished with 

irregular oval shape. After receiving, the PtRu-solid, XRu = 0.3 disk was polished 

28

with emery paper (400 and 600 grit size) and then mirror polished with 9.0, 3.0, 

1.0, 0.3 and 0.05 µm alumina powder (Buehler) to get a smooth crystal faces. 

The Kel-F rod (1.25 cm diameter, Boedeker Plastics, Shiner, TX) was used as a 

supporting material for sealing the disks for fabricating electrodes. Intended for 

proper fastening, one end of Kel-F rod needed to be drilled according to the 

exact shape of the metal disks. For the asymmetric oval shaped disks (PtRu- 

solid, XRu = 0.1 and 0.3), a complicated procedure was followed to seal the edges 

of the electrodes for the present work. In the first step, the PtRu-solid, XRu = 0.1 

and 0.3 disks were scanned with a high resolution (600 dpi) scanner to get an 

identical model picture (Figure 2.1). From these scanned models, exact 

sketches that specify precise geometric dimensions of the peripheral part of the 

uneven disks were drawn with the help of a special 3D computer program 

(AUTODESK INVENTOR 1). This computer program efficiently scaled the 

original photographs of the disks. Then the CNC (COMPUTER NUMERICAL 

CONTROL) milling machine was programmed with scaled dimension to cut the 

Kel-F rod. A hollow cavity that matched the original profile of irregular disk 

resulted. The disks were wrapped with Teflon tape (Zeus, Orangeburg, SC) to 

overcome any mismatch between the drilled Kel-F rod and the disk exterior 

structure. Each metal disk (either regular or irregular shaped) was pressure 

sealed at the edges into the end of drilled Kel-F rod with the polished face 

exposed. In addition, a thin film of silicone sealant was applied to the disk edges 

to avoid further leaking of electrolyte solution onto the backside of the electrode. 

Then the electrodes were soaked in a stirred solution of 0.1M H2SO4 overnight to 

29

leach out the organic compounds which can possibly deteriorate the 

performances of the electrodes. The opposite end of rod was drilled to permit 

positioning of a Pt wire against the backside of the electrode disk. A film of silver 

epoxy (Dynaloy, Indianapolis, IN) secured the contact between the wire and disk. 

The Pt and Au disk working electrodes were periodically polished 

mechanically with alumina from 1.0 µm to down to 0.05 µm followed by 

sonication to remove debris. The PtRu surfaces were polished gently with 0.05 

µm alumina, cleaned by sonication in 5 M KOH, dipped briefly in 5 M H2SO4, and 

finally rinsed in ultra pure water prior to electrochemical measurements [7]. 

Electrochemistry 

Reported cyclic voltammograms of the bulk PtRu and nanometer scale 

catalyst materials were recorded using a conventional glass electrochemical cell. 

The working and counter electrodes and reversible hydrogen electrode (RHE) 

reference were held in the same compartment and immersed in approximately 20 

mL of 0.1 M H2SO4. 

Electrolysis experiments were carried out using a electrochemical cell 

(Figure 2.2) that accommodated a small volume (50 µL) of sample [8]. The 

details of the cell and cell performances will be described in Chapter V. Briefly, a 

spectroelectrochemical cell was employed with a window that enabled the 

formation of a 50 µL thin layer. The window was constructed from a MACOR 

glass ceramic disk with the dimensions 2.54 cm diameter by 0.635 cm thickness 

(Accuratus, Phillipsburg, NJ). A cavity of 50 µL volume was milled into the center 

of one face. A 2 mm recess cut into the top edge of the well allowed the working 

30

Table 2.1: Sample composition in atomic % based on X-ray fluorescence and Xray 

Diffraction [5]. 

Pt PtRu (XRu = 0.1) PtRu (XRu = 0.3) PtRu (XRu = 0.9) 

Ru % 0.0 9.7 29.8 90.5 

Structure fcc fcc fcc hcp 

a (Å) 3.9231 3.9166 3.8907 2.7178 

fcc-face centered cubic; hcp-hexagonal close packed. 

31

(b) 

(a) 

Figure 2.1: Scanned photographs of bulk PtRu disks. (a) XRu= 0.1, (b) XRu= 0.3. 

32

electrode to be positioned reproducibly against the window with the metal disk 

facing the well and in contact with solution [8]. The electrode geometric area 

exposed to solution during electrocatalysis was approximately 0.74 cm 2 . A hole 

drilled through the disk center accommodated the transfer of solution into a gas 

tight syringe (Hamilton Company, Reno, Nevada) through a 23 gauge needle. 

The counter electrode consisted of a ring of Pt-wire and was positioned in the 

center of the cell behind the working electrode. The reference electrode was 

held in a separate compartment behind a wetted stopcock that connected to the 

main body of the cell through a segment of glass tubing. A KCl saturated silver- 

silver chloride electrode (Ag/AgCl) was typically employed as the reference for 

electrolysis experiments. Potentials measured versus the Ag/AgCl reference 

were converted to the RHE scale and reported as volts versus RHE (VRHE). 

The response of the electrolysis cell (Figure 2.2) used for this work was 

characterized by performing cyclic voltammetry measurements using hexamine 

ruthenium (III) trichloride as a probe. Cyclic voltammograms of Ru(NH3)6Cl3 in 

0.1 M NaNO3 were recorded in a conventional electrochemical cell and in the thin 

layer cell. A polycrystalline gold electrode (0.481 cm 2 ) was used as the working 

electrode for these experiments. The Au was pressure sealed into the end of 

Kel-F rod (Boedeker Plastics, Shiner, TX). The voltammetric measurements 

were performed by cycling between -0.3 and 0.4 VAg/AgCl with varying 

concentrations and scan rates. Cell performance to pure polycrystalline Pt 

electrode material (0.81 cm 2 ) were checked by recording cyclic voltammograms 

between 0.0 to 1.5 VRHE at different scan rates and changing the position of the 

33

Cell Body 

Working Electrode 

34 

Clamp 

50 L Cavity 

µ 

Gas Tight 

Syringe 

Macor Window 

Figure 2.2: Front portion of micro-volume electrochemical cell used for sample 

electrolysis.

electrode with respect to cell window (MACOR) position. The cell 

characterization will be described thoroughly in Chapter V. 

Prior to all electrochemical experiments, the solution in the 

electrochemical cell was degassed by bubbling with Ar for ca. 30 min. During 

measurements, Ar flowed in the space above the solution continuously. The 

procedure for sample electrolysis involved first degassing the 0.1M H2SO4 

electrolyte with the working electrode pulled away from the window, followed by 

cycling the electrode between 0.05 VRHE - 0.75 VRHE to verify cleanliness. The 

working electrode was then inserted into the cavity with the potential held at 0.05 

VRHE. Using a gas tight syringe, 25 µL of solution was withdrawn from the cavity 

and replaced with an equal volume of solution containing 2.0 M CH3OH in 0.1 M 

H2SO4 to bring the CH3OH concentration in the cavity to 1.0 M. Following a 5 s 

equilibration period, the electrode was stepped to the reaction potential and held 

for 180 sec. At the end of the electrolysis period, 50 µL of solution was 

withdrawn from the cavity and diluted (by a factor 2-5) in 0.1 M H2SO4 to bring 

the H2CO concentration into the range of 2-25 µM. 

Electrochemical cell potentials were maintained with a three electrode 

potentiostat (PC4/300, Gamry Instruments, Warminster, PA). 

For all bulk electrodes, surface cleanliness was checked by observing 

expected features in the hydrogen adsorption and oxide regions of a cyclic 

voltammogram recorded at 50 mV/s between 0.05 VRHE and 1.5 VRHE for Pt-solid; 

0.05 VRHE and 0.75 VRHE for PtRu alloys; 0.05 VRHE and 1.8 VRHE for Au 

electrode. 

35

Experiments with catalyst were performed by depositing a thin film of 

material on a cleaned, polished and optically flat Au electrode according to 

literature procedures [9, 10]. Catalyst was suspended in ultra pure water (or THF 

for SC-PtRu material) to a known concentration of 2-4 mg/mL and dispensed with 

a micro-pipette onto the Au disk. If needed, the suspension (for the catalysts 

dispersed in water ) was evenly dispersed on the surface by adding additional 

water droplets and then the electrode surface was allowed to air dry for 

overnight. The catalyst films were allowed to air dry for overnight. For all cases 

the resulting electrode surface contained a thin catalyst layer. To remove 

surface contaminants, the electrode was cycled briefly between -0.2 VRHE and 

+1.5 VRHE (for C/Pt, 10 wt%; Pt- black), or 0.0 VRHE and 0.9 VRHE (for PtRu 

catalysts) until the voltammetry reached a steady state and waves characteristic 

of Pt and Ru appeared. The sonochemically prepared materials were activated 

prior to cycling to remove THF derived adsorbates by stepping the potential to 

1.4 VRHE (PtRu, XRu = 0.1) or 1.1 VRHE (PtRu, XRu = 0.5) for 6 s. 

Formaldehyde Detection 

Dissolved H2CO was determined with a fluorescence assay [11]. 

Following electrolysis, sample (50 µL) was reacted with 1, 3- cyclohexanedione 

in an ammonia/ammonium acetate buffer using a flow injection arrangement 

(Figure 2.3) [12, 13]. The reagent solution was prepared by mixing 25 mg of 

CHD, 16.5 g of ammonium acetate and 6.5 mL concentrated HCl (12 M) in 200 

mL of deionized water and then diluted to 250 mL. Reproducible results were 

obtained by allowing the reagent to stand overnight. As discussed earlier [12], 

36

time is required to attain keto/enol tautomeric equilibrium is attained. After 

injection, the sample (50 µL) and reagent were incubated for 3 minutes at 95 ºC 

and then transported to the flow- through fluorometer. The fluorometer was 

constructed around a liquid-core waveguide [14] utilizing Teflon-AF 2400 tubing 

(Biogeneral Fiber technology, 0.040" o.d, 0.034" i.d.). The sample was excited 

by a UV lamp (BF350-UV1, JKL components, Pacoima, CA) with maximum 

intensity at 365 nm. The fluorescence emission was collected at 460 nm through 

a 30 nm band-pass filter on the photodetector (Intro, Socorro, NM) [14]. 

The calibration curve for formaldehyde was recorded with freshly prepared 

standard solutions in the range of 2-25 µM. The linear calibration plot in Figure 

2.4 demonstrates the assay sensitivity and precision. The limit of detection as 

near 100 nM, which is an order of magnitude more sensitive than the lowest 

formaldehyde level, measured in methanol oxidation experiments. Figure 2.5 

shows the fluorescence signal for standard H2CO solution. 

37

CHD/water 

Reagent Pump 

Sample 

Injection Valve 

To Waste 

Thermostatted 

Reaction Coil 

Teflon AF 

(waveguide) 

PD 

Water Bath 

Optical Fiber 

Sample in 

38 

UV Lamp 

(365 nm) 

Flow-through 

Fluorometer 

Waste 

Figure 2.3: Block diagram of the automated H2CO analyzer. 

Flow 

Restrictor

Peak Current (nA) 

Figure 2.4: Calibration curve for standard 

H2CO solution. Concentration in the range of 

2x10 -6 -2.5x10 -5 M. 

Peak Current (nA) 

5 

4 

3 

2 

1 

0 

8 

6 

4 

2 

0 

0 

y = 171.91x + 111.09, R 2 = 0.9987 

0 10 20 30 

Concentration (µM) 

200 

Time (Seconds) 

39 

400 

2 µM 

5 µM 

10 µM 

20 µM 

25 µM 

Figure 2.5: Fluorescence signal of standard H2CO solution.

References 

1. Korzeniewski, C.; Basnayake, R.; Vijayaraghavan, G.; Li, Z.; Xu, S.; 

Casadonte, D. J., Jr. Surf. Sci. 2004, 573, 100. 

2. Kotranarou, A.; Mills, G.; Hoffman, M. R. Environmental Science 

Technology 1992, 26, 2420. 

3. Mason, T. J., In Practical Sonochemistry- Users Guide to Applications in 

Chemistry and Chemical Engineering; Ellis Horwood Publishers, Ltd: New 

York, 1991, pp 46. 

4. Suslick, K. S.; Casadonte, D. J. J. Am. Chem. Soc. 1987, 109, 3459. 


Chem. 1993, 97, 12020. 


Electrochem. Soc. 1994, 141, 1795. 


Weaver, M. J. Electrochim. Acta 1994, 40, 91. 

8. Huang, H.; Korzeniewski, C.; Vijayaraghavan, G. Electrochim. Acta 2002, 

42, 3675. 

9. Friedrich, K. A.; Henglein, F.; Stimming, U.; Unkauf, W. Colloids Surf. A.: 

Physiochem. Eng. Asp. 1998, 134, 193. 

10. Solla-Gullon, J.; Montiel, V.; Aldaz, A.; Clavilier, J. J. Electroanal. Chem. 

2000, 491, 69. 


12. Fan, Q.; Dasgupta, P. K. Anal. Chem. 1994, 66, 551. 

13. Dong, S.; Dasgupta, P. K. Environ. Sci. Technol. 1987, 21, 581. 

14. Dasgupta, P. K.; Genfa, Z.; Li, J.; Boring, C. B.; Jambunathan, S.; Al-Horr, 

R. Anal. Chem. 1999, 71, 1400. 

40

CHAPTER III 

METHANOL ELECTROCHEMICAL CONVERSION TO FORMALDEHYDE 

OVER BULK METAL ELECTRODES 

Introduction 

Parallel reactions (Figure 1.4) during methanol electrooxidation at bulk 

metal electrodes lead to the formation of considerable amounts of H2CO together 

with carbon dioxide (CO2) [1, 2]. In this reports, H2CO yields from methanol 

electrochemical oxidation over smooth Pt and PtRu electrode materials at 

potentials in the range of 0.5-0.8 VRHE. The H2CO generated from 1.0 M 

methanol in 0.1 M H2SO4 at fixed potentials was determined with a small volume 

electrolysis arrangement following short electrolysis periods (180s) using a 

sensitive fluorescence assay [3, 4]. 

With Pt electrodes, the H2CO yields approached 78% at 0.5 VRHE and 

decreased to 13% with increasing potentials (0.8 VRHE). At present, there is a 

general consensus that PtRu alloy materials are the most promising among all 

the binary catalysts for methanol oxidation. Exploring the bulk PtRu alloy 

materials is very important as this material is novel in the area of methanol 

electrochemical conversion to principal product (CO2) or side products (H2CO, 

formic acid, methyl formate, etc). The study demonstrates that the H2CO yield is 

lower (41%, 13% and 39% for XRu = 0.1, 0.3 and 0.9, respectively at 0.5 

VRHE) for PtRu electrodes compared to the Pt electrode. On bulk PtRu 

electrodes, the H2CO yields reached lowest values over the potential range 

41

studied (0.8 V – 0.5 VRHE) for the XRu = 0.3 alloy. The enhancement effect of Ru 

on Pt catalysts for the anodic methanol oxidation can be explained through 

bifunctional mechanism and already being described in Chapter I. The higher 

H2CO yield in bulk electrode materials relative to porous nanometer scale 

catalysts (Chapter IV) are believed to arise from the diffusion of partial oxidation 

products to the solution from smooth metal surfaces without further oxidization of 

H2CO to CO2. 

Experimental 

The polycrystalline Pt and PtRu bulk metal electrodes with three different 

Ru composition (XRu = 0.1, 0.3 and 0.9) were employed to investigate the percent 

yield of H2CO from methanol electrochemical oxidation. Just prior to the 

experiment, all the electrodes were mechanically and electrochemically polished 

and alloys were cleaned according to the procedures described in Chapter II. A 

KClsat Ag/AgCl electrode was employed for electrolysis experiment and a 

reversible hydrogen electrode was used for cyclic voltammetric measurementts. 

Short period (6s) conditioning of PtRu alloy electrodes were done at 1.4 VRHE 

until features appear in hydrogen adsorption desorption region. Cyclic 

voltammetric measurements were performed by cycling between 0.05 to 0.75 

VRHE and 0.05 to 1.5 VRHE for PtRu and Pt electrodes, respectively. The scan 

rate for cyclic voltammetry measurements was 50 mV/s. Electrolysis 

experiments were carried out for 180 s in a 50 µL, argon purged, electrochemical 

cell (Figure 2.2). The electrode was held at 0.0 VRHE during methanol injection. 

The studied potential range was 0.5-0.8 V RHE. Immediately after electrolysis, 50 

42

µL of solution was removed from the cell and diluted (dilution factor 2-5) with 0.1 

M H2SO4 to bring the H2CO concentration into the range of 2-25 µM. Other 

experimental details are described in Chapter II. 

Formaldehyde Determination 

H2CO was determined fluorometrically following reaction with 1, 3- 

cyclohexadione (see Chapter II) using the flow injection arrangement shown in 

Figure 2.3. 

Results and Discussion 

Figures 3.1 and 3.2 show cyclic voltammograms of the PtRu and Pt bulk 

electrode materials, respectively, recorded in a conventional three electrode 

electrochemical cell. The voltammograms recorded in 0.1M H2SO4 demonstrate 

the background response of the materials in aqueous acid solution and the 

cleanliness of the cell, electrolyte and electrode surfaces. Figure 3.1 also 

represents the comparative cyclic voltammograms of Pt (Figure 3.1a) and PtRu 

with XRu = 0.1, 0.3 and 0.9 (Figure 3.1 b, c, d). The positive potential was 

restricted to 0.75 VRHE to minimize oxidation-induced changes in the surface Ru 

composition of PtRu materials [5]. In cyclic voltammetry studies, mainly 

highlighted are (i) the hydrogen region, which contains the hydrogen adsorption 

and hydrogen desorption peaks in the potential range 0.0-0.3 VRHE, and (ii) 

double layer region in the potential range of 0.3-0.7 VRHE. These potential 

regions are important and the most meaningful in the context of cyclic 

voltammetry of pure Pt [6]. Therefore, the voltammogram (Figure 3.1 a) recorded 

43

with a polycrystalline Pt electrode can be considered as a model to describe 

features in these key potential regions. Sharp peaks due to hydrogen adsorption 

(Pt + H + + e - →Pt-H) and hydrogen desorption (Pt-H → Pt + H + + e - ) are observed 

between ~ 0.05 and 0.35 VRHE and are a characteristic of the presence of Pt. 

The double layer charging region occurs between about 0.35 and 0.7 VRHE and is 

narrow for Pt. Figure 3.2 shows a cyclic voltammogram of Pt when the positive 

potential was extended to 1.5 VRHE. In addition to the hydrogen region and 

double layer charging region, features typical for polycrystalline Pt appear in the 

(0.0 to 0.4 VRHE) oxide potential region (0.8 to 1.4 VRHE). Here, an oxide stripping 

peak was observed on the negative sweep at about 0.75 VRHE, and at about 1.0 

VRHE on the forward scan, a shoulder wave associated with Pt oxidation (Pt + 

H2O → PtOH + H + + e - ) appears. 

For clarity, the above mentioned potential regions will be referred to in the 

discussion of cyclic voltammetry of Pt-Ru alloys. Cyclic voltammograms for PtRu 

alloy electrode materials are shown in Figure 3.1 (b), (c), (d). The responses 

demonstrate the effect of increasing the bulk and surface composition of Ru in 

the PtRu alloy electrodes. The cyclic voltammetry results enabled comparisons 

of the hydrogen regions and the double layer regions for the bulk alloy electrodes 

and the pure Pt electrode. While the Pt electrode displayed the typical hydrogen 

adsorption and desorption peaks, as well-defined surface oxidation and reduction 

peaks (Figure 3.1 a and Figure 3.2), the intermetallic alloy electrodes typically 

showed very little, if any, evidence of hydrogen adsorption and desorption 

depending on the Pt/Ru ratio. For all PtRu alloy electrodes, the hydrogen regions 

44

(at ~ 0.05 to 0.35 VRHE) and double layer regions (0.35 to 0.7 VRHE) were 

observed in the same scan range as the polycrystalline Pt electrode. The 

addition of Ru is accompanied by major changes, which include an increase in 

the double layer capacitance, a decrease in the hydrogen adsorption peaks and 

a change in the oxide stripping peak. The PtRu alloy with XRu = 0.1 (Figure 3.1b) 

displayed features that are most similar to those of polycrystalline Pt among the 

alloy samples studied [6, 7]. The incorporation of a small amount of Ru (XRu = 

0.1) caused the hydrogen adsorption and hydrogen desorption peaks to 

diminished slightly and the double layer region to broaden compared to bulk Pt 

electrode. Increasing the Ru amount (XRu = 0.3) (Figure 3.1c), lead to a larger 

double layer thickness and a significant decline in the sharpness of the hydrogen 

adsorption area. The voltammetry of the alloy with XRu = 0.9 composition (Figure 

3.1d) does not display distinct peaks in the hydrogen adsorption region. It is 

instead dominated by a large capacitive current that extends into the double- 

layer region. These findings are consistent with observations for bulk PtRu 

electrodes with different Ru surface compositions (XRu = 0.07, 0.33 and 0.46) and 

have been attributed to the presence of ruthenium oxide on the electrode 

surface, and the adsorption of oxygen like species extending progressively 

toward more negative potentials with increasing Ru composition[6-8]. This 

phenomenon plays a significant role in the conversion efficiency of methanol to 

H2CO, which will be addressed in discussions that follow. 

All results are summarized in Tables 3.1 and Figure 3.3 to Figure 3.6. 

Table 3.1 lists the charge passed, the H2CO quantity produced and the resulting 

45

H2CO yield for the experiments. Below 0.5 VRHE, the % yield value could not be 

determined because the reaction current, and thus the amount of charge passed, 

was too low to be measured. Results for the studied potentials are explained 

below. Special emphasis will be on the lowest positive potentials (0.6 and 0.5 

VRHE) studied as one of the goals of this work was to investigate the reaction 

pathways at low methanol oxidation over potentials. The computation of percent 

conversion was performed by taking the ratio of the charge required to produce 

the quantity of H2CO measured with the fluorescence technique to the total 

charge passed during electrolysis and multiplying the result by 100. In general, 

the H2CO is one of the most dominant by-products of methanol oxidation for 

reactions on bulk metal electrodes with smooth surfaces [9-13]. 

Our study with bulk electrode materials (Table 3.1) shows that the H2CO 

yield generally decreases as the reaction potential is stepped positive. Figure 

3.3, 3.4, 3.5 and Figure 3.6 show the graphical representation of the H2CO yield 

with respect to potential for the Pt-solid and the PtRu-solid with XRu = 0.1, 0.3 

and 0.9 bulk metal electrodes respectively. The topmost entries in Table 3.1 are 

for the Pt- solid electrode (see also Figure 3.3), which showed the highest H2CO 

yields for the considered working potentials (0.8 to 0.5 VRHE) than those of the 

PtRu- solid metal electrodes (2 nd , 3 rd and 4 th entries of Table 3.1). For pure Pt, 

the lowest H2CO yield (13%) was observed at 0.8 VRHE. The formation of H2CO 

over Pt gradually increased toward lower electrode potentials up to 0.5 V where it 

attained 78% of the total electrolysis charge. The H2CO yield started to become 

most significant at potentials lower than 0.7 VRHE. This general behavior was 

46

Current 

0 

0 

0 

0 

0.0 

0.02 mA / cm 2 

E / V vs. RHE 

47 

(a) Pt - solid 

(b) PtRu - solid (XRu = 0.1) 

(c) PtRu - solid (XRu = 0.3) 

(d) PtRu - solid (XRu = 0.9) 

0.2 

0.4 

0.6 0.8 

Figure 3.1: Cyclic voltammograms of bulk electrode materials employed in the 

study. Scans were recorded in Ar purged 0.1 M H2SO4. Polycrystalline Pt 

electrode (Pt-solid, a); bulk PtRu alloys (PtRu-solid) with Ru mole fraction (XRu) of 

0.1 (b), 0.3 (c) and 0.9 (d). Cyclic voltammetric measurements were performed 

by scanning the potential between 0.05 to 0.75 VRHE. The scan rate was 50 mV/s.

Current 

Polycrystalline Pt-solid Electrode 

in 0.1 M H2SO4 

0.05 mA / cm 2 

0 

0.0 0.4 0.8 1.2 

E / V vs. RHE 

Figure 3.2: Cyclic voltammogram of a polycrystalline Pt electrode in Ar purged 

0.1 M H2SO4 recorded at a scan rate of 50 mV/s. 

48 

1.6

observed previously by our research group [9, 14]. These studies employed a 

lower methanol concentration, but the maximum H2CO formation was near 30% 

at 0.5 VRHE and dropped to less than 10% at higher potentials [9, 14]. Using 

HPLC, which is similar to fluorescence in its ability to quantify H2CO in this 

application, Iwasita and co-workers measured percent yield of 71% and 81% for 

H2CO produced at 0.6 VRHE at polycrystalline Pt and Pt(111) electrodes, 

respectively [2]. Furthermore, using the DEMS (Differential Electrochemical 

Mass Spectrometry) technique Wang and co-workers observed H2CO yields of 

up to 72% in reactions at 0.75 VRHE on a smooth polycrystalline Pt electrode in 

0.5 M H2SO4 containing 0.01M methanol [10, 13]. Our results demonstrate, the 

percent conversion of methanol to H2CO on a polycrystalline Pt electrode 

increases as the reaction potential decreases between 0.8V and 0.5 VRHE. The 

result is consistent with conclusions drawn by the Wieckowski group recently 

[15], in a study that confirmed methanol decomposition on Pt proceeds effectively 

through the formation of soluble products at low potentials and that methanol 

dehydrogenation to H2CO can accounts for a majority of the oxidation charge. 

Moreover, the decreasing yield of H2CO observed in our study at potentials 

greater than 0.5 VRHE (Table 3.1 and Figure 3.3) is attributable to the facile water 

activation which takes place at these potentials on Pt and can accelerate CO2 

formation from methanol [16]. 

In contrast to Pt materials, PtRu bulk alloys show evidence of much lower 

secondary product (H2CO) formation, particularly below 0.6 VRHE. The superior 

catalytic performance of PtRu binary alloy materials relative to Pt is well known 

49

and it is explained by the term “bi-functional mechanism” [6, 17, 18]. This term 

put emphasis on the joint activities of both metals, Pt being the one adsorbing 

methanol and cleaving the C-H bond and Ru, the one that promotes the 

dissociation of water and oxidizes the adsorbed residues at low potentials 

compared to pure Pt. Similar to Pt- solid, the PtRu - solid (XRu = 0.1) electrode 

(Figure 3.4 and 2 nd entries of Table 3.1) displays an increase in H2CO yield in 

progressing from higher to lower electrode potentials (from 0.8 VRHE to 0.5 VRHE). 

In the case of the lowest working potential (0.5 VRHE), the H2CO yield (41%) is 

reduced substantially from that of the bulk platinum electrode (78%). A very low 

yield of H2CO (~6.2%) and higher reaction charge were found for the highest 

reaction potential (0.8 VRHE). At 0.6 VRHE, relatively lower yield (32%) of H2CO 

was achieved if compared to the yield on pure Pt (38%) but the reaction yield 

remains almost the same for both electrodes at 0.7 VRHE. The results indicate 

that incorporating a small amount of Ru into Pt enhances the catalytic activity of 

the surface with respect to the oxidation of methanol beyond H2CO. This result 

will be discussed in greater detail later in the chapter. 

Of all the bulk alloys, PtRu solid (XRu = 0.3) showed the best catalytic 

performance and lowest H2CO yields (3 rd entries of Table 3.1 and Figure 3.5) for 

all studied working potentials (0.8 to 0.5 VRHE). On the lowest Ru content alloy 

(XRu = 0.1), the methanol surface chemistry was similar to that of polycrystalline 

Pt, and there is only a small change of H2CO yield above 0.5 VRHE. With the 

inclusion of a medium amount of Ru (XRu = 0.3), drastic changes were observed 

in reaction charges, H2CO quantity and H2CO yields. Comparing the yield (see 

50

Figure 3.7) for the lowest two working potentials (0.6 and 0.5 VRHE), PtRu solid 

with XRu = 0.3 produced appreciably lower yields (8.4% and 13% for 0.6 and 0.5 

VRHE, respectively) than those of PtRu solid with XRu = 0.1 (32% and 41% for 0.6 

and 0.5 VRHE, respectively) and pure Pt (38% and 78% for 0.6 and 0.5 VRHE, 

respectively). Here the role of Ru-adatoms is accelerating the methanol 

oxidation rate on smooth Pt at low potentials and the high coverage of oxides on 

Ru sites is expected to enable the conversion of H2CO to products with greater 

oxygen content, such as HCOOH and CO2. Also, the PtRu-solid (XRu = 0.3) 

electrode used in our study showed excellent catalytic activity in the H2CO yield 

(~3-5%) for higher electrode potentials (0.8 and 0.7 VRHE). The very low yields 

for higher electrode potentials are due to the combined effect of higher driving 

force for the reaction and enhancement of water activation by Ru atoms. 

The measurement with PtRu-solid, XRu = 0.9 showed higher average 

H2CO yields (4 th entries of Table 3.1 and Figure 3.6) compared to that of PtRu- 

solid, XRu = 0.3 electrode for all potentials (0.8 to 0.5 VRHE). The H2CO yields 

were significantly higher at 0.6 and 0.5 VRHE for PtRu-solid, XRu = 0.9 (16% and 

39% respectively) relative to PtRu-solid, XRu = 0.3 (8.4% and 13% respectively). 

The lower activity of the PtRu solid with XRu = 0.9 towards methanol oxidation 

relative to the PtRu solid with XRu = 0.3 can be explained by the lower availability 

of Pt atoms for methanol dehydrogenation, which requires an ensemble, namely, 

a specific number of Pt sites for methanol decomposition [7, 19]. The alloy with 

higher Ru mole fraction contains a passivating surface layer, probably Ru oxides, 

which can inhibit methanol dissociative chemisorption at ambient temperature 

51

[20]. The jump in H2CO yield measured in reactions over PtRu, XRu = 0.9 at 0.5 

V may reflect a decrease in the rate of surface oxide attack on carbon containing 

fragments at the low electrochemical driving force together with an inhibition of 

C-H bond breaking steps by the oxides. These effects would create a trap for 

accumulation of H2CO by reducing the fraction of reacting CH3OH molecules that 

progress from H2CO to HCOOH, or from H2CO to CO and subsequently CO2. 

These findings indicate that the bulk alloys with the high Pt or Ru composition is 

not as active towards the methanol oxidation as the alloy containing a medium 

(XRu = 0.3) amount of Ru. 

The data in Table 3.1 indicate that the PtRu solid with small amount of Ru 

(XRu = 0.1) behaves almost like pure Pt with respect to H2CO yields, whereas 

incorporating somewhat more Ru (XRu = 0.3), improves the catalytic activity of 

the alloy toward methanol oxidation significantly. In literature reports, at ambient 

temperatures PtRu electrodes with low Ru surface concentration (XRu = 0.1- 0.3) 

show the best kinetics for the electrooxidation of methanol (ranging from 0.005 to 

0.5 M) [6, 7]. Our PtRu - solid (XRu = 0.1) is highly active at the most positive 

potentials and also showed better performances for lowest electrode potentials 

regarding the H2CO yield compared to pure Pt electrode. The current results 

indicate that the PtRu solid with XRu = 0.3 is the best performing electrode among 

those studied for 1.0 M methanol oxidation due to the small H2CO yields shown 

at low reaction potentials in combination with high reaction currents. The Pt/Ru 

ratio in alloy materials is not the only factor for attaining the best catalytic activity. 

Reactant concentration also plays a role. In earlier work by our research group, 

52

Table 3.1: Summary of charge and formaldehyde yields from the oxidation of 1.0 

M CH3OH on bulk Pt and PtRu electrode materials a . 

Electrode Materials b Potential c 

(E/V vs. 

RHE) 

Average 

Charge 

(mC) 

53 

Average H2CO 

Quantity (nmol) 

H2CO 

Yield % 

Pt-solid 0.8 3.08 2.12 13 ± 4 

0.7 4.33 4.5 20 ± 3 

0.6 5.69 10.7 38 ± 7 

0.5 0.33 1.33 78 ± 4 

PtRu – solid (XRu = 0.1) 0.8 8.63 2.89 6.2 ± 3.1 

0.7 2.39 2.82 23 ± 3 

0.6 1.77 3.14 32 ± 7 

0.5 0.41 0.85 41 ± 8 

PtRu – solid (XRu = 0.3) 0.8 25.95 4.65 3.5±0.1 

0.7 10.3 1.69 3.8±1.8 

0.6 14.1 3.86 8.4 ± 5.3 

0.5 3.86 2.10 13 ± 4 

PtRu – solid (XRu = 0.9) 0.8 4.11 1.27 6.5 ± 2.1 

0.7 7.02 2.06 5.8 ± 0.7 

0.6 3.20 2.43 16 ± 5 

0.5 1.51 3.02 39 ± 2 

a All experiments involved the electrolysis of 1.0 M CH3OH in 0.1 M H2SO4 in a volume of 50 µL for a period 

of 180s. The values reported for the charge and H2CO amount are averages calculated from results of 3-5 

trials. The yield values were computed from the ratio of the charge required to produce H2CO to the total 

charge passed during the electrolysis period. These yield values were averaged to give the reported 

percentages and uncertainties. The uncertainties represent 95.5% confidence limits based on two times the 

sample standard deviation. b The following electrode materials were used: Polycrystalline Pt electrode (Ptsolid); 

bulk PtRu alloys (PtRu-solid) with Ru mole fraction (XRu) of 0.1, 0.3 and 0.9. c The reference electrode 

was KCl saturated Ag/AgCl. For reporting purposes, all the measured potentials vs. Ag/AgCl were converted 

to the reversible hydrogen electrode (RHE) scale and expressed as volts versus RHE (VRHE).

% H 2CO yield 

100 

90 

80 

70 

60 

50 

40 

30 

20 

10 

77.85 

37.67 

54 

19.81 

12.61 

0 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 

Figure 3.3: H2CO yields from methanol electrochemical oxidation on a 

polycrystalline Pt electrode. All experiments involved the electrolysis of 1.0 M 

CH3OH in 0.1 M H2SO4 in a volume of 50 µL for a period of 180s.

% H 2CO yield 

100 

90 

80 

70 

60 

50 

40 

30 

20 

10 

41.42 

32.32 

55 

23.09 

0 

0.45 0.5 0.55 0.6 0.65 

E/V vs. RHE 

0.7 0.75 0.8 0.85 

Figure 3.4: H2CO yields from methanol electrochemical oxidation on a PtRu- 

solid (XRu = 0.1) electrode. All experiments involved the electrolysis of 1.0 M 


6.21

% H 2CO yield 

100 

90 

80 

70 

60 

50 

40 

30 

20 

10 

13 

8.41 

0 

3.8 

3.52 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 




56

% H 2CO yield 

100 

90 

80 

70 

60 

50 

40 

30 

20 

10 

38.73 

15.65 

0 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 




57 

5.8 

6.5

% H 2CO yield 

90 

80 

70 

60 

50 

40 

30 

20 

10 

0 

Pt-solid PtRu-solid (XRu=0.1) PtRu-solid (XRu=0.3) PtRu-solid (XRu=0.9) 

Electrode Composition 

58 

0.6 V vs.RHE 

0.5 V vs.RHE 

Figure 3.7: Comparison of H2CO yields from methanol (1.0 M) electrochemical 

oxidation on bulk electrode materials at 0.6 and 0.5 VRHE.

it was shown that the surface coverage of adsorbed residues from methanol 

decomposition, especially CO, appears to increase on Pt surfaces with 

increasing methanol concentration. Poisoning by the adsorbed species is 

assumed to be responsible for limiting CO2 formation below about 0.8 VRHE [21]. 

In addition, the Ru composition influences the balance between CO2 and H2CO 

yields from methanol and Iwasita and co-workers [22] reported that an improved 

catalytic effect can be obtained if the final oxidation step ((CO)ad + 

(OH)ad→CO2+H + +e - ) occurs between CO adsorbed at Pt and OH adsorbed at Ru 

(Pt(CO)ad + Ru(OH)ad→CO2+H + +e - ) [22]. In our studies of H2CO formation at the 

PtRu (XRu = 0.1) alloy, 10% Ru atoms seems inadequate to provide sufficient 

hydroxide for oxidizing the high amount of adsorbed residues expected. Our 

results are in good agreement with the work of Iwasita, et al, which showed that 

the rate of methanol oxidation to CO2 (Figure 1.4) is faster at XRu = 0.5 than at 

XRu = 0.15 PtRu alloy [23, 24] even though the XRu = 0.15 surface gave higher 

current densities for methanol oxidation. These workers concluded that the lower 

Ru content alloy produced a greater yield of H2CO, consistent with our findings. 

With our techniques for electrolysis and H2CO determinations, it is 

possible to make repetitive measurements simply and rapidly. Figure 3.8 and 3.9 

summarize results of the analysis of data for polycrystalline Pt from a series of 

repetitive electrolysis experiments at fixed potentials. Figure 3.8 shows the plot 

of nanomoles of H2CO formed versus the reaction charge passed following the 

sample electrolysis. The Y-axis of the plot in Figure 3.8 gives the H2CO quantity 

determined from the intensity of H2CO bands in comparison to a calibration curve 

59

(see Figure 2.4). At each potential, the error bars represent 95% confidence 

intervals on the mean of 3 or 4 repetitive measurements. Figure 3.9 reports the 

H2CO formation rate as a function of potential. The ability to make such 

determinations enables the significance of reaction rate changes due to the 

different electrode properties (i.e., composition, surface structure) to be assessed 

statistically, and uncertainties to be assigned to kinetic and thermodynamic 

parameters derived from the data. In the same manner, this type of data 

analysis is summarized for the PtRu alloy electrodes in Figure 3.10 and Figure 

3.11 (XRu = 0.1); Figure 3.12 and Figure 3.13 (XRu = 0.3); and Figure 3.14 and 

Figure 3.15 (XRu = 0.9). 

In general Figure 3.8 (Pt electrode), Figure 3.10 (PtRu-solid, XRu = 0.1 

electrode), Figure 3.12 (PtRu-solid, XRu = 0.3) and Figure 3.14 (PtRu-solid, XRu = 

0.9) compare uncertainty in H2CO determinations to the error in corresponding 

coulometry measurements. The errors relative to the averages are larger for the 

coulometry than the fluorescence spectroscopy measurements. Under the 

conditions of the experiments, the precision of the fluorescence technique 

outperforms that of the electrochemistry. 

From the graphs of the rate of H2CO formation versus electrode potential 

shown in Figure 3.9 (Pt electrode), Figure 3.11 (PtRu-solid, XRu = 0.1 electrode), 

Figure 3.13 (PtRu-solid, XRu = 0.3) and Figure 3.15 (PtRu-solid, XRu = 0.9), the 

general trend is that the rate of formation of H2CO is highest at the lower 

electrode potentials, 0.5 and 0.6 VRHE. 

60

Conclusions 

Although the in situ IR and DEMS methods are highly efficient in the study 

of competing methanol electrooxidation pathways, they are incapable of the 

direct detection of H2CO. As we [9, 14] and others [2, 10, 13] have shown, H2CO 

can be generated at low potentials in reactions on bulk electrode surfaces and 

should be considered as an important intermediate in mechanistic work. The 

reported fluorescence assay had sufficient sensitivity to track H2CO produced in 

a micro volume cell employing polycrystalline Pt and bulk PtRu alloy working 

electrodes over short electrolysis periods (180 s). The catalytic activity of PtRu 

electrodes toward methanol oxidation was superior to polycrystalline Pt, but the 

yields of H2CO were strongly dependent on the Ru content. The lowest H2CO 

yield was attained for the alloy electrode with average Ru composition (XRu = 0.3) 

over the potential range studied (0.8 V – 0.5 VRHE). The PtRu- solid (XRu = 0.3) 

electrode is considered to be in the range of optimum composition for the 

electrochemical efficient oxidation of methanol in PEM type fuel cells. 

61

Formaldehyde quantity(nmol) 

12 

10 

8 

6 

4 

2 

0.5 V RHE 

0 

0 1 2 3 4 5 6 7 

charge (mC) 

62 

0.8 V RHE 

0.7 V RHE 

0.6 V RHE 

Figure 3.8: Plot of the nanomoles H2CO formed vs. the reaction charge passed 

following the oxidation of 1.0 M methanol on a polycrystalline Pt electrode in 0.1 

M H2SO4 (Reaction time = 180 s).

Rate (nmol/s/cm 2 ) 

0.08 

0.07 

0.06 

0.05 

0.04 

0.03 

0.02 

0.01 

0 

0.4 0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 

Figure 3.9: Plot of H2CO formation rate averaged over 180 s electrolysis periods 

vs. potential for the oxidation of 1.0 M methanol in 0.1 M H2SO4 on a 

polycrystalline Pt electrode. 

63

Formaldehyde quantity (nmol) 

3.5 

3 

2.5 

2 

1.5 

1 

0.5 

0.5 V RHE 

0.6 V RHE 

0.7 V RHE 

0 

0 1 2 3 4 5 

Charge (mC) 

6 7 8 9 10 

64 

0.8 V RHE 


following the oxidation of 1.0 M methanol on a bulk alloy PtRu- solid (XRu = 0.1) 

electrode in 0.1 M H2SO4 (Reaction time = 180 s).


0.025 

0.02 

0.015 

0.01 

0.005 

0 

0.4 0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 


vs. potential for the oxidation of 1.0 M methanol in 0.1 M H2SO4 on a bulk alloy 

PtRu- solid (XRu = 0.1) electrode. 

65

Formaldehyde Quantity (nmol) 

6 

5 

4 

3 

2 

1 

0.5 V RHE 

0.7 V RHE 

0.6 V RHE 

0 

0 5 10 15 

Charge (mC) 

20 25 30 

66 

0.8 V RHE 





0.04 

0.035 

0.03 

0.025 

0.02 

0.015 

0.01 

0.005 

0 

0.4 0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs RHE 




67

Formaldehyde quantity (nmol) 

3.5 

3 

2.5 

2 

1.5 

1 

0.5 

0.5 V RHE 

0.6 V RHE 

0 

0 1 2 3 4 

Charge (mC) 

5 6 7 8 

68 

0.8 V RHE 

0.7 V RHE 





0.03 

0.025 

0.02 

0.015 

0.01 

0.005 

0 

0.45 0.5 0.55 0.6 0.65 

E/V vs. RHE 

0.7 0.75 0.8 0.85 




69

References 

1. Cao, D.; Lu, G.-Q.; Wieckowski, A.; Wasileski, S. A.; Neurock, M. J. 

Phys.Chem. B. 2005 , 109 , 11622. 

2. Batista, E. A.; Malpass, G. R. P.; Motheo, A. J.; Iwasita, T. J. Electroanal. 

Chem. 2004 , 571 , 273. 

3. Dong, S.; Dasgupta, P. K. Environ. Sci. Technol. 1987 , 21 , 581. 

4. Nash, T. Biochem. J. 1953 , 55 , 416. 


Weaver, M. J. Electrochim. Acta 1994 , 40 , 91. 


Chem. 1993 , 97 , 12020. 


Chem. 1994 , 98 , 617. 

8. Gasteiger, H. A.; Markovic, N. M.; Ross, P. N., Jr. J. Phys. Chem. 1995 , 

99 , 16757. 

9. Childers, C. L.; Huang, H.; Korzeniewski, C. Langmuir 1999 , 15 , 786. 

10. Wang, H.; Wingender, C.; Baltruschat, H.; Lopez, M.; Reetz, M. T. J. 

Electroanal. Chem. 2001 , 509 , 163. 

11. Jusys, Z.; Behm, R. J. J. Phys. Chem. B 2001 , 105 , 10874. 

12. Jusys, Z.; Kaiser, J.; Behm, R. J. Langmuir 2003 , 19 , 6759. 

13. Wang, H.; Loffler, T.; Baltruschat, H. J. Appl. Electrochem. 2001 , 31 , 

759. 

14. Korzeniewski, C.; Childers, C. L. J. Phys. Chem. B 1998 , 102 , 489. 

15. Lu, G.-Q.; Chrzanowski, W.; Wieckowski, A. J. Phys. Chem. B 2000 , 104 

, 5566. 

16. Corrigan, D. S.; Weaver, M. J. J. Electroanal. Chem. 1988 , 241 , 143. 

17. Watanabe, M.; Motoo, S. J. Electroanal. Chem. 1975 , 60 , 267. 

70

18. Hamnett, A., In Interfacial Electrochemistry. Theory, Experiment, and 

Applications ; Wieckowski, A., Ed.; Marcel Dekker: New York, 1999, pp 

843. 

19. Jarvi, T. D.; Stuve, E. M., In Electrocatalysis ; Lipkowski, J. and Ross, P. 



Electrochem. Soc. 1994 , 141 , 1795. 

21. Kardash, D.; Korzeniewski, C. Langmuir 2000 , 16 , 8419. 

22. Iwasita, T. Electrochim. Acta 2002 , 47 , 3663. 

23. Iwasita, T.; Hoster, H.; John-Anacker, A.; Lin, W. F.; Vielstich, W. 

Langmuir 2000 , 16 , 522. 

24. Batista, E. A.; Hoster, H.; Iwasita, T. J. Electroanal. Chem. 2003 , 554-555 

, 265. 

71

CHAPTER IV 

METHANOL ELECTROCHEMICAL CONVERSION TO FORMALDEHYDE 

OVER SUPPORTED CATALYST MATERIALS 

Introduction 

The DMFC is considered to have a great deal of potential due to its 

simplified system design and direct use of liquid fuel. But the present 

performance levels are not yet sufficient for commercial applications. The limited 

performance of the cell is due in part to poor kinetics of the anode reaction. 

Improvements in methanol oxidation catalysts are desired [1, 2]. Pt-based 

materials are typically employed. On pure Pt, methanol oxidation is strongly 

inhibited by poison formation. Bimetallic catalysts, such as PtRu, have much 

greater activity and resist poisoning [3, 4]. 

In contrast to bulk metals, the electrodes in fuel cell are comprised of 

nanometer-scale metal particles adsorbed onto either micron-scale carbon 

particles or porous carbon supports [2]. The catalytic, physical, electrochemical 

and electronic properties of nanostructure catalyst materials can be quite 

different from those of bulk electrode materials [2, 5]. Reaction rates have been 

shown to be sensitive to the size and structure of the nanoparticle catalyst [6], 

and phase separation can occur in bi-metallic nanoparticles [7, 8]. In the study of 

methanol electrochemical conversion to H2CO over bulk electrodes (Chapter III), 

a significant amount of H2CO was produced. For comparison purposes, an 

attempt was made to study the methanol electrochemical oxidation reaction 

72

pathways over nanometer scale catalyst materials, which have more practical 

applications than bulk electrodes. In contrast to the solid metal electrodes 

studied, the yields of H2CO over catalyst particles were low (below 10%) during 

the electrochemical oxidation of 1.0 M methanol in 0.1M H2SO4 for short 

electrolysis periods (180 s). Materials examined include: Pt- black; C/Pt, 10 wt%; 

JM PtRu, XRu = 0.5; and SC PtRu catalyst in the following Ru compositions XRu = 

0.1, 0.25, and 0.5. In fuel cells, PtRu materials with XRu= 0.5 are favored for 

DMFC operations [2]. However, in nanometer scale PtRu catalysts, phase 

separation can occur to produce Ru-rich and Pt-rich regions [8, 9]. This structure 

contrasts bulk PtRu, which form random alloys. In addition, the high surface area 

particles create a porous bed that enables reacting molecules to undergo 

encounters with multiple catalyst sites and thereby facilitate more complete 

oxidation of reactant. This chapter presents the yield data for H2CO produced in 

reactions of methanol over nanometer-scale catalyst materials in the MACOR 

based small volume (50 µL) electrochemical cell described earlier (Chapter II). 

The oxidation of 1.0 M methanol at potentials in the range of 0.5-0.8 VRHE was 

examined. 

Experimental 

The details of the catalysts preparation and the electrochemical techniques 

employed were described in Chapter II. 

Formaldehyde Determination 

H2CO was determined fluorometrically following reaction with 1, 3- 

73

cyclohexadione (see Chapter II) using the flow injection arrangement shown in 

Figure 2.3. 


Cyclic voltammograms of a polycrystalline gold (Au) electrode before 

(Figure 4.1) and after modification with an ultra-thin layer of Pt (Figure 4.2) or 

PtRu based catalyst materials (Figure 4.3) are displayed. The voltammograms 

were recorded in a conventional three electrode electrochemical cell. The cyclic 

voltammogram in Figure 4.1 for a mirror polished Au electrode in Ar purged 0.1 

M H2SO4 between H2 and O2 evolution potentials provides information about the 

cleanliness of surface. The shape and peak potentials for the waves that appear 

are in good agreement with prior reports [10]. The oxide formation peaks on the 

positive going sweep between 1.2 and 1.4 VRHE (associated with Au + H2O → 

AuOH + H + + e - ) and the oxide stripping peak (AuOH + H + + e - → Au + H2O) on 

the negative going sweep near 1.0 VRHE, together with the featureless double 

layer region (0.0 to 0.85 VRHE) represent the basic features of a clear, 

polycrystalline gold electrode. There are no obvious faradic electrochemical 

processes occurring in the wide potential region between 0.0-1.0 VRHE. This 

potential region provides a broad window over which the characteristics features, 

for instance, hydrogen adsorption and desorption peaks (0.0-0.3 VRHE) and 

double layer charging (0.3-0.7 VRHE) from adsorbed Pt and PtRu nanoparticles 

can be observed. 

Figure 4.2 shows the cyclic voltammogram of the Au electrode after 

adsorption of carbon supported platinum catalyst (C/Pt, 10 wt %) (Figure 4.2a) or 

74

a Pt-black film (Figure 4.2b) in 0.1 M H2SO4. The faradic current relative to the 

charging current in each case suggests the catalyst film coverage is a few 

monolayers [11, 12]. The cyclic voltammograms of the C/Pt, 10 wt% and Pt- 

black film (Figure 4.2 a and b) have waves characteristic of platinum similar to 

Figure 3.2 in Chapter III. In acid electrolyte solutions, features associated with 

hydrogen adsorption and oxide formation and stripping are present, but are 

considerably weaker for C/Pt, 10 wt% than for bulk polycrystalline Pt. The current 

and peak potentials for C/ Pt, 10 wt% catalyst are consistent with prior reported 

voltammograms of the catalyst recorded on Au in 0.05 M H2SO4 [12, 13]. In 

contrast to the carbon supported Pt material, which are affected by the 

capacitance of the carbon, particles of Pt-black adsorbed on Au (Figure 4.2 b) 

display waves in the clear 0.1 M H2SO4 electrolyte solution that are more typical 

of bulk polycrystalline Pt. For Pt-black, the defined features in the hydrogen 

adsorption and oxide formation and strippinng regions are analogous to those 

observed for Pt-black catalyst immobilized Au or carbon [14,15, 16] and 

moderate to high metal loading (≥ 20%) carbon supported Pt catalysts on Au [11- 

13, 17]. The large oxide stripping peaks were found on the negative going 

sweep near 0.7 VRHE and 0.65 VRHE for Pt-black and C/Pt ,10 wt% respectively. 

For both cases the hydrogen adsorption/desorption region in between 0.0 to 0.3 

VRHE indicate the presence of Pt on the gold surface. On the positive scan, the 

current rise near 0.8 VRHE and the Pt-oxide formation features are observed in 

the range of 0.9 to 1.5 VRHE. The clear definitions of all the waves indicate the Pt 

particles have a high degree of cleanliness. The arrow in both cyclic 

75

Current 

Polycrystalline Au electrode 

in 0.1 M H SO 

2 4 

0 

0.02 mA / cm 2 

0.0 0.4 0.8 1.2 1.6 2.0 

E / V vs. RHE 

Figure 4.1: Cyclic voltammogram of a polycrystalline gold electrode in 0.1 M 

H2SO4; Scan rate: 50 mV/s. 

76

Current 

Au / catalyst film 


200 µA 

0 

0 

-0.1 0.5 1.1 1.7 

E / V vs. RHE 

77 

(a) C/Pt, 10wt% 

(b) Pt- black 

Figure 4.2: Cyclic voltammogram of a polycrystalline gold electrode in 0.1 M 

H2SO4 after adsorption of C/Pt, 10 wt% (a) or Pt-black (b) catalyst films; Scan 

Rate: 50 mV/s. Arrow represents the stripping wave for gold oxide.

voltammograms represents the stripping wave for gold oxide, which indicates 

that gold is exposed at the surface between Pt particles. 

Figure 4.3 shows the representative cyclic voltammograms of a Au 

electrode after modification with an ultra thin layer of nanometer scale PtRu 

catalyst materials (JM PtRu, XRu = 0.5; SC PtRu, XRu = 0.1 and SC PtRu, XRu = 

0.5) used in this study. The response of the catalyst particles can be correlated 

to the Ru content in the particles by comparison to results for bulk PtRu alloy 

electrodes (Chapter III). For the sonochemically prepared nanoparticles with XRu 

= 0.1 in 0.1 M H2SO4 (Figure 4.3b, central voltammogram), waves appear at 

0.05-0.3 VRHE characteristic of hydrogen adsorption and hydrogen desorption [11, 

12]. The features for the SC PtRu, XRu = 0.1 nanocatalysts are almost similar to 

those of Pt electrode (Figure 3.2) except some lessening of wave current in the 

hydrogen region reflecting the presence of Ru. Like the bulk PtRu alloy electrode 

with XRu = 0.07 studied earlier [18, 19], the SC PtRu, XRu = 0.1 electrode also 

displayed a much narrower double layer charging region compared to the Ru rich 

electrode. Our PtRu bulk electrode (Figure 3.1) and catalyst (Figure 4.3 b) with 

similar Ru content (XRu = 0.1) display the same properties, which is encouraging. 

The upper and bottom voltammograms in Figure 4.3 represent the commercial 

JM PtRu, XRu = 0.5 (Figure 4.3a) and SC PtRu, XRu = 0.5 (Figure 4.3c) 

nanoparticle materials, where the multiple hydrogen adsorption and desorption 

peaks are completely suppressed. In these cases, incorporating more Ru (XRu = 

0.5) into the particles, increases the Ru surface composition and the hydrogen 

78

Current 

Au / catalyst film 


0 

0 

0 

200 µA 

(a) JM PtRu, XRu = 0.5 

(b) SC PtRu, XRu = 0.1 

(c) SC PtRu, XRu = 0.5 

0.0 0.2 0.4 0.6 0.8 

E / V vs. RHE 

Figure 4.3. Cyclic voltammograms of catalyst materials employed in the study. 

Scans were recorded in Ar purged 0.1 M H2SO4 at 50 mV/s. The following 

nanoscale catalyst materials were studied as thin films adsorbed on a Au 

electrode- Johnson Matthey (JM) PtRu with XRu = 0.5 (a); PtRu catalyst 

prepared via a sonochemical (SC) method with XRu = 0.1 (b) and 0.5 (c). 

79 

X 2 

X 4

adsorption and desorption features become less resolved and overlap 

increasingly with the double layer region. The broad envelopes typical for high 

Ru content materials are noticeable in the range of 0.3-0.7 VRHE, the double layer 

charging region. This section is featureless typical of the pseudocapacitive 

properties of Ru oxide [9]. The same findings have been reported by other 

researchers where the formation of surface oxide occurs at progressively lower 

potentials as the Ru coverage on a pure Pt catalyst was increased [18, 20-22]. 

The H2CO produced with these and a few additional catalysts on a gold 

electrode was determined as a function of potential (0.5-0.8 VRHE). Table 4.1 and 

4.2 report reaction charge and the quantity of H2CO produced during short 

periods (180s) of methanol electrochemical oxidation over the nanoscale 

catalysts. Figure 4.4 to 4.11 graphically represent the H2CO yields with respect 

to electrode potentials for all the catalysts studied. The percent conversion was 

computed from the ratio of the charge required to produce the quantity of H2CO 

detected fluorometrically to the total charge passed during electrolysis. In Table 

4.1, the results for Pt based catalysts (Pt-black and C/Pt, 10 wt %) are 

summarized. For Pt-black the greatest yield (5.28%) was found at 0.6 VRHE and 

decreased at higher potentials. The largest yield for C/Pt, 10 wt% catalyst was 

7.9% for 0.5 VRHE and decreased continuously to 0.5 % with increasing 

potentials. High reaction charges were observed for these catalysts at almost all 

potentials compared to a solid platinum electrode (Chapter III, Table 3.1). The 

response reflects both the larger surface area of the catalyst and the strong 

poisoning that occurred for the solid platinum at the potentials employed. Similar 

80

Table 4.1 Summary of charge and formaldehyde yields from the oxidation of 1.0 

M CH3OH over Pt based nanoscale catalyst materials a . 


(E/V vs. 

RHE) 

Average 

Charge 

(mC) 

81 

Average 

H2CO 

Quantity 

(nmol) 

H2CO 

Yield % 

Pt- Black 0.8 121.21 13.00 2.0 ± 0.5 

0.7 64.30 10.68 3.3 ± 0.5 

0.6 35.08 9.76 5.3 ± 0.7 

0.5 17.01 3.70 4.8 ± 0.5 

C/Pt, 10 wt% 0.8 121.21 3.39 0.5 ± 0.1 

0.7 65.15 7.57 2.2 ± 0.3 

0.6 33.76 3.95 2.3 ± 0.2 

0.5 14.78 5.70 7.9 ± 0.6 

a All experiments involved the electrolysis of 1.0 M CH3OH in 0.1 M H2SO4 in a volume of 50 µL 

for a period of 180s. The values reported for the charge and H2CO amount are averages 

calculated from results of 3-5 trials. The yield values were computed from the ratio of the charge 

required to produce H2CO to the total charge passed during the electrolysis period. These yield 

values were averaged to give the reported percentages and uncertainties. The uncertainties 

represent 95.5% confidence limits based on two times the sample standard deviation. b The 

following catalyst materials were used on a gold electrode: Pt-Black (from JM); C/Pt, 10 wt% = 

Pt at 10 wt% on Vulcan XC-72R carbon (from JM). c The reference electrode was KCl saturated 

Ag/AgCl. For reporting purposes, all the measured potentials vs. Ag/AgCl were converted to the 

reversible hydrogen electrode (RHE) scale and expressed as volts versus RHE, (VRHE).

% H 2CO Yield 

6 

5 

4 

3 

2 

1 

0 

4.762 

82 

5.281 

3.251 

1.973 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs. RHE 

Figure 4.4: H2CO yields from methanol electrochemical oxidation in 0.1 M H2SO4 

on Pt -Black catalyst on a polycrystalline Au electrode. All experiments involved 

the electrolysis of 1.0 M CH3OH in 0.1 M H2SO4 in a volume of 50 µL for a period 

of 180s. 

% H 2CO Yield 

10 

8 

6 

4 

2 

0 

7.85 

2.29 

2.24 

0.54 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 


on C/Pt, 10 wt% catalyst on a polycrystalline Au electrode. All experiments 

involved the electrolysis of 1.0 M CH3OH in 0.1 M H2SO4 in a volume of 50 µL for 

a period of 180s.

esults have been observed with lower methanol concentration and different 

catalyst materials [23, 24]. With the higher methanol concentrations employed in 

this study, the yield of partial oxidation products is expected to increase as the 

decrease in interfacial water activity lowers the rate of surface oxide formation 

(H2O → OHads + H + + e - ) [25-27]. 

To enable comparisons among the catalyst systems, JM PtRu, XRu = 0.5, 

and sonochemically prepared PtRu catalysts with various Ru surface 

compositions (XRu = 0.1, 0.25 and 0.5), were chosen. Similar to PtRu bulk 

electrodes, PtRu catalysts have considerably higher tolerance for carbon 

monoxide than pure Pt materials [18]. The PtRu catalysts are predicted to have 

a higher catalytic activity for methanol oxidation than platinum due to the 

bifunctional mechanism (See Chapter III) [3, 4, 18]. In addition, incorporating Ru 

into the Pt lattice is believed to weaken the Pt-COads bond strength, reducing the 

–COads coverage [28]. Compared to Pt catalyst, the PtRu materials showed 

superior activity toward CH3OH oxidation, and the average yields for H2CO are 

lower than for the reactions on PtRu materials particularly at 0.5 VRHE, except for 

SC PtRu, XRu = 0.1 (Table 4.2 and Figure 4.7- 4.11). Similar to the bulk 

materials, the H2CO yield for the SC PtRu, XRu = 0.1 catalyst more closely 

resembles that for clean Pt, as they both produce a larger quantity of H2CO (yield 

~7.8%) at the lowest working potentials (0.5VRHE). Although all catalysts in our 

analysis generate very small yields, and there is no sharp difference in yield 

values for all catalysts across the studied potential range (0.5-0.8 VRHE), it is 

interesting to note that at 0.5 VRHE the H2CO yield declined from 7.4% to 0.5% 

83

with increasing Ru content (Figure 4.11). PtRu catalysts with XRu = 0.5 showed 

the lowest H2CO yields, and JM PtRu, XRu = 0.5 and SC PtRu, XRu = 0.5 results 

are comparable to each other. At the lowest and the most important electrode 

potential (0.5 VRHE) from the standpoint of practical fuel cell applications, the 

H2CO yield did not decrease significantly with the incorporation of a small amount 

of Ru (~7.4% yield for XRu = 0.1) compared to Pt based catalysts (~8% yield for 

C/Pt, 10 wt%). However, the H2CO yield decreased with further increases in Ru 

to 1.4% for SC PtRu, XRu = 0.25 and 0.5-0.7% for both SC PtRu, XRu = 0.5 and 

JM PtRu, XRu = 0.5. Similar trends were observed in this project for bulk 

electrodes (Chapter III) and for nanometer scale catalyst by other researchers 

[20, 29]. DEMS results for PtRu catalysts showed that the alloy materials 

containing nearly equal Pt and Ru content (XRu = 0.3 to 0.5) exhibited higher 

activity toward CO2 production from methanol at ~ 0.4- 0.5 VRHE compared to 

lower Ru containing catalysts [20]. The observed result demonstrated that the 

surface area normalized methanol oxidation activity of PtRu catalysts and 

electrodes is highest for low Ru contents around XRu = 0.15 at more positive 

potentials (~ 0.7 VRHE) [20]. These activities refer to the inherent chemical 

activity obtained by normalizing the oxidation current to the active surface area 

determined by COad- stripping to CO2 [20]. The DEMS study with smooth Pt and 

PtRu (XRu = 0.35 - 0.45) materials by Wang and co-workers showed that the 

current efficiency of CO2 increased with the incorporation of Ru as it promotes 

the methanol oxidation on Pt at low potentials [29]. At room temperature, PtRu 

alloy catalysts having a Ru composition XRu= 0.3-0.5, at present the best 

84

Table 4.2: Summary of charge and formaldehyde yields from the oxidation of 1.0 

M CH3OH over PtRu based nanoscale catalyst materials a . 


Average H2CO H2CO Catalyst 

(E/V vs. Charge Quantity Yield % loading 

RHE) (mC) (nmol) 

JM PtRu, XRu = 0.5 0.8 28.28 3.00 2.2 ± 0.6 62µg/cm 2 

0.7 24.06 2.99 1.6 ± 0.5 

0.6 8.81 0.39 1.1 ± 0.7 

0.5 5.58 0.16 0.5 ± 0.1 

SC PtRu, XRu = 0.5 0.8 27.47 0.59 0.7 ± 0.5 86µg/cm 2 

0.7 18.87 0.51 2.0 ± 0.7 

0.6 15.83 0.45 2.1 ± 0.6 

0.5 14.92 0.22 0.7 ± 0.6 


0.7 34.92 7.29 2.8 ± 0.3 

0.6 8.17 2.34 5.5 ± 0.1 

0.5 7.28 0.47 1.4 ± 0.5 


SC PtRu, XRu = 0.1 

with Nafion 

0.7 37.00 6.81 3.5 ± 0.4 

0.6 19.04 4.78 5.2 ± 0.6 

0.5 7.43 2.79 7.4 ± 0.7 

0.8 153.17 5.25 0.7 ± 0.3 96 µg/cm 2 

0.7 118.50 6.34 1.1 ± 0.5 

0.6 84.57 4.40 1.0 ± 0.1 

0.5 18.72 1.76 2.0 ± 0.5 

b Johnson Matthey (JM) PtRu with Ru mole fraction (XRu) 0.5; PtRu catalyst prepared via a sonochemical 

(SC) method with XRu = 0.1, 0.25 and 0.5. SC PtRu, XRu = 0.1 with Nafion refers to the catalyst film capped 

by a Nafion coating. a,c Other conditions are the same with table 4.1 

85

% H 2CO yield 

10 

8 

6 

4 

2 

7.38 

5.19 

86 

3.53 

4.61 

0 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 


on SC PtRu, XRu= 0.1 catalyst on a polycrystalline Au electrode. All experiments 


a period of 180s. 

% H 2CO yield 

4 

3 

2 

1 

1.95 

1.004 

1.11 

0.71 

0 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 


on SC PtRu, XRu= 0.1 catalyst, capped by Nafion on a polycrystalline Au 

electrode. All experiments involved the electrolysis of 1.0 M CH3OH in 0.1 M 

H2SO4 in a volume of 50 µL for a period of 180s.

% H 2CO yield 

6 

4 

2 

0 

1.36 

5.52 

87 

2.8 

4.43 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 




a period of 180s. 

% H 2CO yield 

5 

3 

1 

-1 

0.66 

2.09 

1.96 

0.67 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 




a period of 180s.

% H 2CO yield 

6 

5 

4 

3 

2 

1 

0 

0.53 

1.11 

88 

1.57 

2.15 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 

Figure 4.10: H2CO yields from methanol electrochemical oxidation in 0.1 M 

H2SO4 on JM PtRu, XRu= 0.5 catalyst on a polycrystalline Au electrode. All 

experiments involved the electrolysis of 1.0 M CH3OH in 0.1 M H2SO4 in a 

volume of 50 µL for a period of 180s.

% H 2CO yield 

9 

8 

7 

6 

5 

4 

3 

2 

1 

0 

Pt-black C/Pt, 10 wt% SC PtRu, 

XRu=0.1 

1 M Methanol; 0.5 V vs. RHE 

SC PtRu, 

XRu=0.1, with 

Nafion 

89 

SC PtRu, 

XRu=0.25 

Electrode Composition 

JM PtRu, 

XRu=0.5 

SC PtRu, 

XRu=0.5 

Figure 4.11: H2CO yield from methanol (1.0 M) electrochemical oxidation on 

catalyst materials at 0.5 VRHE.

catalysis for CO2 production from methanol, probably due to the high surface 

coverage of oxides on Ru. The bottom entry of Table 5.2 indicates that the 

methanol oxidation occurs more efficiently when catalyst (SC PtRu, XRu = 0.1) 

was immobilized in Nafion. For every reaction potentials, the H2CO yield for SC 

PtRu, XRu = 0.1, with Nafion, became much less than that of SC PtRu, XRu = 0.1, 

without Nafion. At the most highlighted investigated potential (0.5 VRHE), the 

H2CO yield and H2CO quantity were reduced to a greater extent as the catalyst 

layer and the products and byproducts produced, are entrapped by Nafion film. 

The reaction charges were also increased a great deal compared to the catalyst 

without a Nafion film covering. 

In the platinum based catalysts (Table 4.1), the amount of H2CO produced 

at 0.5 VRHE are 3.70 nmol (Pt-black) and 5.70 nmol (C/Pt, 10 wt %) which were 

somewhat large relative to the amount (~ 0.1-0.4 nmol) in the PtRu based 

catalysts with average or higher Ru composition (XRu = 0.25 -0.5). This 

difference in H2CO quantity formed probably arises from the Ru incorporation 

and availability of oxides to transform H2CO to higher oxidation products. The 

more ruthenium is alsol catalyzing the CO oxidation, thus favoring the CO2 

formation rather than forming much H2CO [3, 4, 18]. A more useful value to 

analyze is the H2CO percent yield because the quantity is corrected for surface 

coverage effects. With both the Pt-black and the Pt-Ru deposits, the percent 

yield is very low which is below 5% for 0.5 VRHE. Since the percent yield values 

are well below the solid Pt and PtRu electrodes (Chapter III), these experiments 

indicate that the small nanoparticles grouping of Pt or PtRu catalysts better drive 

90

the reaction beyond H2CO. In the nano-scale catalyst films partial oxidation 

product like H2CO generally desorbs from catalyst surfaces and diffuses within 

the film where it can meet additional catalyst for more complete oxidation of 

methanol [24]. 

The analysis of data and the results for repetitive measurements are 

summarized graphically in Figure 4.12 to Figure 4.25 for all the studied catalysts. 

Figure 4.12 to Figure 4.18 report H2CO formation rate as a function of potential. 

Figure 4.19 to Figure 4.25 shows the plot of nanomoles vs. the reaction charge 

passed following the sample electrolysis. The error bars represent 95% 

confidence intervals on the mean of 3 or 4 repetitive measurements. 

Conclusions 

Pt and PtRu catalyst materials were supported as a thin layer on a Au 

electrode and characterized by cyclic voltammetry in 0.1 M H2SO4 and for activity 

toward H2CO production from 1.0 M CH3OH. The studies showed good precision 

and reproducibility. All catalysts produced high electrolysis currents, but low (< 

10%) yields of H2CO, indicating CH3OH and any intermediate H2CO produced is 

converted to more complete oxidation products over the catalyst bed. In this 

respect, the catalyst layers are much more reactive toward CH3OH than bulk 

electrodes of similar composition. Among the materials studied, the PtRu 

catalysts with XRu = 0.5 showed the lowest yield for H2CO. It appears that both 

the physical structure of catalyst materials (bulk versus nanometer scale 

particles) and composition influence CH3OH oxidation pathways. In this work, 

91

the responses observed for bulk and nanometer scale materials are compared 

and discussed further in Chapter VI. 

92


0.045 

0.03 

0.015 

0 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs. RHE 


vs. potential for the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a Pt -Black 

catalyst layer adsorbed on a polycrystalline Au electrode. 


0.03 

0.02 

0.01 

0 

0.4 0.5 0.6 0.7 0.8 0.9 

E/V vs RHE 


vs. potential for the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a C/Pt, 10 

wt% catalyst layer adsorbed on a polycrystalline Au electrode. 

93


0.035 

0.03 

0.025 

0.02 

0.015 

0.01 

0.005 

0 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 


vs. potential for the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a SC PtRu, 

XRu= 0.1 catalyst layer adsorbed on a polycrystalline Au electrode. 


0.01 

0.008 

0.006 

0.004 

0.002 

0 

0.4 0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 



XRu= 0.1 catalyst layer (capped by Nafion) adsorbed on a polycrystalline Au 

electrode. 

94


0.025 

0.02 

0.015 

0.01 

0.005 

0 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 





0.0007 

0.0006 

0.0005 

0.0004 

0.0003 

0.0002 

0.0001 

0 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 

E/V vs. RHE 




95


0.007 

0.006 

0.005 

0.004 

0.003 

0.002 

0.001 

0 

0.45 0.5 0.55 0.6 0.65 0.7 0.75 0.8 0.85 0.9 

E/V vs. RHE 


vs. potential for the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a JM PtRu, 



16 

14 

12 

10 

8 

6 

4 

2 

0.6 V RHE 

0.5 V RHE 

0.7 V RHE 

0 

0 20 40 60 80 100 120 140 

Charge (mC) 

96 

0.8V RHE 


following the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a Pt-Black catalyst 

layer adsorbed on a polycrystalline Au electrode (Reaction time = 180 s).


10 

8 

6 

4 

2 

0 

0.5 V RHE 

0.6 V RHE 

0.7 V RHE 

0 20 40 60 80 100 120 140 

Charge (mC) 

97 

0.8 V RHE 


following the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a C/Pt, 10 wt% 

catalyst layer adsorbed on a polycrystalline Au electrode (Reaction time = 180 s). 


16 

14 

12 

10 

8 

6 

4 

2 

0 

0.5 V RHE 

0.6 V RHE 

0.7 V RHE 

0 10 20 30 40 50 60 70 

Charge (mC) 


following the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a SC PtRu, XRu= 

0.10 catalyst layer adsorbed on a polycrystalline Au electrode (Reaction time = 

180 s). 

0.8 V RHE


7 

6 

5 

4 

3 

2 

1 

0.5 V RHE 

0 

0 20 40 60 80 100 120 140 160 180 

Charge (mC) 

98 

0.7 V RHE 

0.6 V RHE 

0.8 V RHE 



0.10 catalyst layer (capped by Nafion) adsorbed on a polycrystalline Au electrode 

(Reaction time = 180 s). 


12 

10 

8 

6 

4 

2 

0 

0.5 V RHE 

0.6 V RHE 

0 5 10 15 20 25 30 35 40 45 50 

Charge (mC) 




180 s). 

0.7 V RHE 

0.8 V RHE


0.7 

0.6 

0.5 

0.4 

0.3 

0.2 

0.1 

0.5 V RHE 

0.6 V RHE 

0.7 V RHE 

0 

12 14 16 18 20 22 24 26 28 

Charge (mC) 

99 

0.8 V RHE 




180 s). 


3.5 

3 

2.5 

2 

1.5 

1 

0.5 

0.5 V RHE 

0.6 V RHE 

0 

0 5 10 15 20 25 30 35 

Charge (mC) 

0.7 V RHE 

0.8 V RHE 


following the oxidation of 1.0 M methanol in 0.1 M H2SO4 over a JM PtRu, XRu= 


180 s).

References 

1. Hogarth, M. P.; Hards, G. A. Platinum Metals Rev. 1996, 40, 150. 

2. Hogarth, M. P.; Ralph, T. R. Platinum Metals Rev. 2002, 46, 146. 

3. Watanabe, M.; Motoo, S. J. Electroanal. Chem. 1975, 60, 267. 

4. Hamnett, A., In Interfacial Electrochemistry. Theory, Experiment, and 

Applications; Wieckowski, A., Ed.; Marcel Dekker: New York, 1999, pp 

843. 

5. El-Sayed, M. A. Accounts of Chemical Research 2001, 34, 257. 

6. Mayrhofer, K. J. J.; Blizanac, B. B.; Arenz, M.; Stamenkovic, V. R.; Ross, 

P. N.; Markovic, N. M. J. Phys. Chem. B 2005 ,109 , 14433. 

7. Stroud, R. M.; Long, J. W.; Swider-Lyons, K. E.; Rolison, D. R. Microsc. 

Microanal. 2002, 8, 50. 

8. Diaz-Morales, R. R.; Liu, R.; Fachini, E.; Chen, G.; Segre, U. C.; Martinez, 

A.; Cabera, C.; Smotkin, E. S. J. Electochem. Soc. 2004, 151, A1314. 

9. Rolison, D. R.; Hagans, P. L.; Swider, K. E.; Long, J. W. Langmuir 1999 , 

15 , 774. 

10. Angnes, L.; Richter, E. M.; Augelli, M. A.; Kume, G. H. Anal. Chem. 2000, 

2000, 5503. 

11. Park, S.; Xie, Y.; Weaver, M. J. Langmuir 2002, 18, 5792. 

12. Park, S.; Tong, Y. Y.; Wieckowski, A.; Weaver, M. J. Electrochemistry 

Communications 2001, 3, 509. 

13. Park, S.; Tong, Y. Y.; Wieckowski, A.; Weaver, M. J. Langmuir 2002 , 18 , 

3233. 

14. Day, J. B.; Vuissoz, P.-A.; Oldfield, E.; Wieckowski, A.; Ansermet, J.-P. J. 

Am. Chem. Soc. 1996, 118, 13046. 

15. Long, J. W.; Rolison, D. R., In New Directions in Electroanalytical 

Chemistry II; Leddy, J., Vanysek, P. and Porter, M. D., Ed.; 

Electrochemical Society: Pennington, NJ, 1999; Vol. PV99-5, pp 125. 

16. Long, J. W.; Ayers, K. E.; Rolison, D. R. J. Electroanal. Chem. 2002, 522, 

58. 

100

17. Park, S.; Wasileski, S. A.; Weaver, M. J. J. Phys. Chem. B 2001 , 105 , 

9719. 


Chem. 1993, 97, 12020. 


Chem. 1994, 98,617. 


21. Long, J. W.; Swider, K. E.; Merzbacher, C. I.; Rolison, D. R. Langmuir 

1999 , 15 , 780. 

22. Lee, C. E.; Bergens, S. H. J. Phys. Chem. B 1998, 102, 193. 

23. Korzeniewski, C.; Childers, C. L. J. Phys. Chem. B 1998, 102, 489. 

24. Childers, C. L.; Huang, H.; Korzeniewski, C. Langmuir 1999 , 15 , 786. 

25. Gao, P.; Chang, I. C.; Zhou, Z.; Weaver, M. J. J. Electroanal. Chem. 

1989, 272,161. 

26. Wasmus, S.; Wang, J.-T.; Savinell, R. F. J. Electrochem. Soc. 1995,142 , 

3825. 

27. Lin, W.-F.; Wang, J.-T.; Savinell, R. F. J. Electrochem. Soc. 1997,144 , 

1917. 



Electroanal. Chem. 2001, 509, 163. 

30. Jarvi, T. D.; Stuve, E. M., In Electrocatalysis ; Lipkowski, J. and Ross, P. 


101

Chapter V 

A MICRO-VOLUME ELECTROCHEMICAL CELL FOR THE STUDY OF 

METHANOL OXIDATION PATHWAYS 

Introduction 

A definition of thin layer electrochemistry is that area of electrochemical 

endeavor in which special advantage is taken of restricting the diffusion field of 

electroactive species and products [1]. It is an approach for achieving bulk 

electrolysis conditions and a large electrode surface area to solution volume ratio 

(A/V) without convective mass transfer and involves decreasing the solution 

volume (V) such that a very small amount is confined to a thin layer adjacent to 

the working electrode surface. The solution thickness (or cell height) should be 

comparable to or smaller than the diffusion layer of analyte under study [2]. In a 

thin-layer flow cell, solution is constrained to a film passing over a planar 

electrode held at a fixed potential. Figure 5.1 illustrates an exploded view of the 

thin layer region. An electrochemically active substance passes over an 

electrode held at a potential sufficiently great (positive or negative) for an 

electron transfer (either oxidation or reduction) to occur. An amperometric 

current is produced that is directly proportional to the concentration of the analyte 

entering the thin layer cell [1]. 

There are many electrochemical studies, including quantification of 

product and by-product formation during methanol electrochemical reaction, 

investigations of adsorption, electrodeposition, complex reaction mechanisms, 

102

and n-value determinations where thin layer cells have been widely used. In 

addition to the benefit of conservation of valuable sample, thin-layer cells are 

also very useful for spectroelectrochemical analysis, because the entire solution 

in the thin-layer can be rapidly and completely electrolyzed and this makes it 

possible to obtain spectroscopic information without interference from unreacted 

analyte in bulk solution [3]. 

In the present studies, a thin layer spectrochemical cell was adapted for 

electrolysis measurements. The conventional flat window was modified to 

accommodate a 50 µL sample volume. In early work a window was constructed 

from Kel-F. However, the permeability of Kel-F to atmospheric O2 potentially 

limits electrolysis measurements at low (≤ 0.5VRHE) potentials. As part of the 

study, a new window materials, MACOR, was tested due to its extremely low 

permeability to gases and easy machinability. A well known redox system, [Ru 

(NH3)6] 2+/3+ was employed to characterize the cell using cyclic voltammetry. The 

cell was used in investigations of methanol electrochemical conversion to H2CO 

as discussed in Chapters III and IV. 

Experimental and Instrumentation 

Reagents 

Sodium Nitrate (NaNO3, Mallinckrodt Chemicals work, Hazelwood, MO) 

and sulfuric acid (H2SO4) solutions were prepared from distilled water that was 

further processed with a 4-cartidge Nanopure II system (Barnstead, Dubuque, 

IA). The Sodium nitrate was reagent grade and used as received. Hexaammine 

ruthenium (III) trichloride (Ru (NH3)6Cl3, Ru 32.6%) was obtained from Alfa Aesar 

103

Figure 5.1: Exploded view of electron transfer at an electrode surface in a thinlayer 

flow cell [1]. 

104

in highest purity and used as received. Solutions were freshly prepared prior to 

measurements. The H2SO4 was obtained in 99.999% purity from Aldrich 

(Milwaukee, WI). Argon (Trinity Gases, Dallas, TX) was of ultrahigh purity. 

Small Volume Electrochemical Cell 

A diagram showing the small volume region of the electrochemical cell 

employed for CH3OH electrolysis was included in Chapter II (Figure 2.2). An 

infrared spectroelectrochemical cell was adapted by substituting a conventional 

window for one that had been modified to accommodate a fixed volume of 50 µL. 

The general design of the cell had been reported earlier [7, 8]. As part of this 

project, improvements were made in reducing the O2 permeability of the cell. 

Figure 5.2 showed the side view and top view of cell window denoting the 

detailed dimensions of each section. The working electrode was positioned 

adjacent to the cavity in a manner that entraps solution. The total well depth is 

2.54 mm. There is a 1.9 mm recess cut into the top edge of the well that helps in 

positioning the working electrode reproducibly with the metal disk facing the well 

and in contact with the solution. The diameter of the outer well and the inner well 

is 12.7 mm and 9.9 mm, respectively. A drilled hole (1.09 mm diameter) through 

the disk center allows the transfer of solution through a 23 gauge needle of a gas 

tight syringe (Hamilton Company, Reno, Nevada). 

Choice of Cell Material and Cell Fabrication 

In studying methanol electrooxidation pathways wheather to by-products 

such as H2CO or the main product CO2, dissolved O2 in the reactant solution can 

105

undergo reduction and thereby cause error in the determination of the methanol 

oxidation charge. In acidic solution, O2 reduction can progress to H2O or H2O2 

depending upon the working electrode material and potential, as shown in Eq. 1: 

O2 + 2 e - + 2 H + ⎯→ H2O2 [1a] 

O2 + 4 e - + 4 H + ⎯→ 2 H2O [1b] 

When O2 reduction occurs simultaneously with CH3OH oxidation, the charge 

necessary for reaction of O2 off-sets the charge produced in the oxidation 

reaction. As a result, the measured charge is lower than the actual charge 

passed in the CH3OH oxidation process. 

One of the objectives of this project was to find a material that is both 

impermeable to O2 and easily machinable. In the cell used earlier [7, 8], a Kel-F 

disk was used as the cell window, because it could be precision machined to 

create the needed 50 µL volume. However, a major limitation of Kel-F in this 

application is that it is permeable to gases and allows diffusion of O2, into the 

cell. The O2 interference becomes problematic as low (below 0.5 VRHE) CH3OH 

oxidation potentials are reached. As an alternative to Kel-F, we initially explored 

the use of glass. Glass windows (25.4 mm diameter by, 6.35 mm thickness) for 

use in UV-visiible spectroscopy cells could be easily obtained commercially. The 

disks were impermeable to O2 and had excellent resistance to attack by 

chemicals. However, they tended to break easily during machining, processes 

which involved grinding the glass by pressing the flat end of a high precision cut 

metal rod against an abrasive powder on the glass surface. A milling machine 

106

well depth = 

0.100 " 

drilled hole = 

0.043 " dia. 

drilled hole = 

0.043 " dia. 

0.5 " 

1.0 " 

Side View 

0.50" 

1.0 " 

Top View 

107 

well depth = 

0.075 " 

0.25 " 

0.39" 

Figure 5.2: Side view and top view of the electrochemical cell window.

was used to carefully align the rod and glass disk during the grinding. 

Further searching brought to light the glass ceramic MACOR as a 

potential material. MACOR is a white, odorless, porcelain-like material 

composed of approximately 55% fluorphlogopite mica (KMg3AlSi3O10F2) in a 

borosilicate glass matrix. It can be melted and cast into various forms using 

conventional glass making techniques. Like Kel-F, MACOR can be easily 

machined. It has extremely low O2 permeability and can be purchased in disk 

forms the size of glass optical windows. Thus, MACOR was tested as a thin 

layer cell material. 


Cell Characterization in 0.1 M H2SO4 

The cell was first tested for cleanliness and response in a conventional 

three electrode electrochemical cell configuration. Cyclic voltammograms were 

recorded with mirror polished polycrystalline Pt electrode in a conventional 

electrochemical cell and then in the electrolysis cell with the electrodes placed at 

different positions relative to the MACOR window. Cyclic voltammograms were 

recorded between -0.2 V and 1.3 V vs. a Ag/AgCl reference electrode and then 

the potentials were converted to the RHE scale. Figure 5.3 compares 

voltammograms obtained at the scan rate of 50 mV/s in a conventional 

electrochemical cell [Figure 5.3 (a)] and the electrolysis cell when the electrode 

was mounted far from MACOR window [Figure 5.3 (b)]. The two voltammograms 

have a similar appearance. The important features are (i) the defined hydrogen 

adsorption and hydrogen desorption peaks in the so called hydrogen region 

108

etween 0.0-0.3 VRHE, (ii) the flat double layer charging region over the potential 

range 0.3-0.7 VRHE, (iii) the peaks associated with Pt oxidation which start at ~ 

0.95 VRHE, and (iv) the Pt reduction peak at ~ 0.65 VRHE. There is a small drop in 

current density for the oxide stripping peak (from 280 µA/cm 2 to 250 µA/cm 2 ) and 

in hydrogen adsorption and desorption regions for the electrode in the 

electrolysis cell. Otherwise, the voltammograms exhibit good agreement with the 

well-known current-voltage response of clean polycrystalline Pt electrodes [9]. It 

can be concluded that when the electrode is positioned in the electrolysis cell far 

from the MACOR window, the cell performance is similar to that of the 

conventional cell. 

In Figure 5.4, cyclic voltammograms recorded using the electrolysis cell 

and a polycrystalline Pt electrode positioned very near to the MACOR window 

are shown. The voltammograms were recorded at two different scan rates: 10 

mV/s [Figure 5.4(a)] and 50 mV/s [Figure 5.4 (b)]. The peak currents become 

lower with decreasing scan rate, as expected, and the position of peaks in the 

oxide stripping (~ 0.6 to 0.8 VRHE) and hydrogen region (~ 0.0 to 0.3 VRHE) does 

not change. 

Figure 5.5 shows comparison voltammograms for the Pt electrode when it 

is positioned outside the MACOR window [Figure 5.5 (a)] and after insertion into 

the MACOR cavity [Figure 5.5 (b)]. The scan rate was slowed to 10 mV/s to 

reduce distortions in the voltammogram that would be caused by resistive effects 

when the electrode is inside the cavity. Looking at the cyclic voltammogram in 

the MACOR window [Figure 5.5 (b)], there is a skewing of the hydrogen 

109

adsorption region (~ at 0.03 VRHE, pointed by gray arrow) which is due to some 

resistance drop. The effect was not observed in the cyclic voltammogram 

recorded with the electrode positioned far from the cavity [Figure 5.4 and Figure 

5.5 (a)]. In both cases [Figure 5.5 (a) and (b)], oxide reduction peaks were well 

resolved and observed at ~ 0.65 VRHE. However, there is a small feature at ~ 

1.1VRHE that could indicate some contamination arising from the MACOR 

material [Figure 5.5 (b)]. This effect was ignored in our study as this 

contamination was not seen in the cell characterization part [Figure 5.6 (c) and 

(d)] and quantitative analysis through electrolysis measurements (Chapter III and 

IV). 

Finally, based on the absence of any significant distortions, the cyclic 

voltammograms presented in Figure 5.3, 5.4 and 5.5 demonstrate that the 

electrochemical measurements can be carried out in this thin layer 

electrochemical cell configuration constructed with MACOR material without 

considerable ohmic drops. 

Cell Characterization using a Reversible Redox Active Probe 

The response of the thin layer cell used in this work was further 

characterized by performing cyclic voltammetry measurements using a soluble, 

reversible redox probe, hexamine ruthenium trichloride [Ru(NH3)6Cl3]. The 

voltammograms were compared to the response of the redox reaction in a 

conventional electrochemical cell operating under the same conditions. The 

working electrode employed for these measurements was a polycrystalline gold 

disk (0.481 cm 2 ) that had been pressure sealed into the end of Kel-F rod 

110

Current 



0.1 mA / cm 2 

0 

0 

0.0 0.4 0.8 1.2 1.6 

E / V vs. RHE 

111 

(a) conventional 

electrochemical 

cell 

(b) Electrolysis Cell 

Figure 5.3: Cyclic voltammograms of a polycrystalline Pt electrode in 0.1 M 

H2SO4 recorded in a conventional electrochemical cell (a) and the electrolysis 

cell with the electrode mounted far from MACOR window (b). The scan rate was 

50 mV/s.

Current 



0.1 mA / cm 2 

0 

0 

(a) Electrolysis Cell,10 mV/s 

(b) Electrolysis Cell, 50 mV/s 

0.0 0.4 0.8 1.2 1.6 

E / V vs. RHE 

112 

X 2.5 


H2SO4 recorded in an electrolysis cell with the electrode mounted near the 

MACOR window. Scan rate was 10 mV/s (a) and 50 mV/s (b).

Current 



2 

0.1 mA / cm 

0 

0 

(a) Electrolysis Cell (Far from 

MACOR window),10mV/s 

(b) Electrolysis Cell (MACOR 

window),10mV/s 

0.0 0.4 0.8 1.2 1.6 

E / V vs. RHE 


H2SO4 recorded in an electrolysis cell with the electrode mounted near the 

MACOR window (a) and mounted inside the MACOR window (Arrow represents 

skewing of the hydrogen adsorption region) (b). The scan rate was 10 mV/s. 

113 

X 3 

X 4

(Boedeker Plastics, Shiner, TX). The voltammetric responses were measured by 

cycling between -0.3 and 0.4 VAg/AgCl and using [Ru(NH3)6Cl3] solution with 0.1 M 

NaNO3 as a supporting electrolyte [Figure 5.6]. Various probe concentrations 

and scan rates were tested. The reaction equation can be written as follows: 

[Ru(NH3)6] 3+ + e - ↔ [Ru(NH3)6] 2+ 

The cyclic voltammograms presented in Figure 5.6 exhibit the characteristic 

reduction peak at ~0.05 VAg/AgCl and oxidation peak at ~ 0.12 VAg/AgCl. In Figure 

5.6 (a) and (b), the voltammograms were recorded in a conventional 

electrochemical cell. At the scan rate of 50 mV/s [Figure 5.6 (a)] with 5.0 x 10 -3 

M [Ru(NH3)6] 3+ , the response appears reversible and limited by linear diffusion. 

Slowing the scan rate to 0.2 mV/s, [Figure 5.6 (b)], gives a response closer to 

that of steady-state. A major concern when working with thin layer cells, 

especially when nonaqueous solutions or very low supporting electrolyte 

concentrations are employed, is the high resistance of solution in the cavity. In 

the measurements undertaken, the reference and counter electrodes were 

located in the bulk solution outside the thin layer and far from the working 

electrode. Thus, the scan rate was slowed to compensate for the slower cell 

response time due to the high resistance. Figure 5.6 (c) shows the 

voltammogram recorded at this slower scan rate (0.2 mV/s) with the electrode 

positioned in the MACOR window. However, the long time required to complete 

the scan and limited quantity of reactant in the cavity (2.5 x 10 -9 nmols / µL x 50 

µL) resulted in analyte depletion during the scan. Thus, the voltammogram in Fig 

5.6c shows the effect of the decreasing [Ru(NH3)6] 3+ . 

114

To overcome the problem of reactant depletion, a higher [Ru(NH3)6] 3+ 

concentration (2.5 x 10 -2 M) and faster scan rates (5-10 mV/s) were tested (data 

not shown). These conditions resulted in a diffusion controlled response 

affected by cell resistance. However, effects of reactant depletion were much 

less, as completing the forward and reverse scans required less time and thus 

less charge was passed. To maintain a balance between cell resistance and 

reactant depletion, scan rates slower than 5-10 mV/s were attempted. A 

reversible voltammetric response that appeared to be limited by linear diffusion 

[Figure 5.6 (d)] was finally established by employing the slower scan rate (1 

mV/s) with 2.5 x 10 -2 M [Ru(NH3)6] 3+ . The voltammogram showed some 

depletion of reactants from the cavity, but very low cell resistance effects. 

Peak separation is a unique characteristic of a thin-layer electrochemical 

cell response in the cyclic voltammetry of a reversible redox couple [1]. In 

theory, the peak potentials for the reduction and oxidation waves should be the 

same in an ideal thin layer cell when the cell boundary falls within the analyte 

diffusion layer. However, in practice, some peak separation occurs due to 

uncompensated solution resistance [10]. In Figure 5.6 (d), the peak separation is 

about 79 mV (and 66 mV in Figure 5.6 (c)), which is close to the theoretical peak 

separation for a cell operating under linear diffusion control (∆Ep = Epa-Epc = 57/n 

mV, where Epa (anodic) and Epc (cathodic) peak potentials) [11]. The result 

indicates that while the thin layer cavity has a small volume, the cell does not 

display true thin layer behavior, since the boundaries are somewhat large 

compared to reactant diffusion layer under the conditions of the measurements 

115

Current 

200 µA/cm 2 

0 

0 

0 

0 

-0.3 -0.2 -0.1 0.0 0.1 0.2 0.3 0.4 

E/V vs. Ag/AgCl 

116 

(a) 50 mV/s 

(b) 0.2 mV/s 

x 5 

(c) 0.2 mV/s 

x 50 

(d) 1 mV/s 

Figure 5.6: Cyclic voltammograms of [Ru(NH3)6Cl3] in 0.1 M NaNO3 recorded in 

a conventional electrochemical cell (a, b) and the MACOR cavity of the thin layer 

electrolysis cell shown in Figure 3.2 (c, d). The [Ru(NH3)6] 3+ concentration and 

scan rates were : (a) 5 × 10 -3 M and 50 mV/s , (b) 5 × 10 -3 M and 0.2 mV/s (c) 2.5 

× 10 -3 M and 0.2 mV/s (d) 2.5 ×1 0 -2 M and 1 mV/s. The scans started from 0.4 

VAg/AgCl and swept toward more negative potentials. 

x 10

Nevertheless, the voltammograms in Figure 5.6 display equal oxidation and 

reduction peak heights, which is a diagnostic for reversibility. Considering peak 

separation and peak height in the cyclic voltammograms, both the conventional 

and the thin layer cells showed comparable results. Additionally, our thin layer 

cell is much better than other recent silicon based designs fabricated through 

micromachining technology. In this latter case, the reported peak separation 

value is 250 mV [12]. 

Conclusions 

This project advanced the design of a small volume electrolysis cell used 

in prior studies by incorporating a machinable MACOR glass ceramic disk 

window, which resists oxygen permeation. The cell response was characterized 

by observing hydrogen adsorption/desorption and oxide formation waves on Pt in 

0.1 M H2SO4 and the reversible, diffusion controlled waves for hexamine 

ruthenium trichloride [Ru(NH3)6Cl3] in cyclic voltammetry measurements. 

117

References 

1. Heineman, W. R.; Kissinger, P. T. Chapter 3, Laboratory Techniques in 

Electroanalytical Chemistry; 2nd Edition ed.; MARCEL DEKKER, INC.: 

New York, 1996. 

2. Bard, A. J.; Faulkner, L. R. Electrochemical Methods; Wiley: New York, 

1980. 

3. Hubbard, A. T.; F.C., A. Electranal. Chem. 1970, 4. 

4. Seki, H.; Kunimatsu, K.; Golden, W. G. Appl. Spectrosc. 1985, 39, 437. 

5. Bethune, D. S.; Luntz, A. C.; Sass, J. K.; Roe, D. K. Surf. Sci. 1988, 197, 

44. 

6. Popenoe, D. D.; Stole, S. M.; Porter, M. D. Appl. Spectrosc. 1992, 46, 79. 

7. Huang, H.; Korzeniewski, C.; Vijayaraghavan, G. Electrochim. Acta 2002, 

42, 3675. 


9. Angerstein-Kozlowska, H.; Conway, B. E.; Sharp, W. B. A. J. Electroanal. 

Chem. 1973, 43, 9. 

10. Tom, G. M.; Hubbard, A. T. Anal. Chem. 1971, 43, 671. 

11. Evans, D. H.; O'Connel, K. M.; Peterson, R. A.; Kelly, M. J. Journal of 

Chemical Education 1983, 60, 290. 

12. Yun, K.-S.; Joo, S.; H-J, K.; Kwak, J.; Yoon, E. Electroanalysis 2005, 17, 

959. 

118

CHAPTER VI 

SUMMARY 

The conversion of methanol to H2CO on Pt and PtRu based electrode 

materials was investigated with the use of electrochemical methods and 

fluorescence spectroscopy. Methanol is a promising fuel for PEM type fuel cells 

and investigation of its electrochemistry over various electrode catalyst materials 

is essential to establish guidelines for DMFC research. The complete oxidation 

of methanol to CO2 is a multi-step process (Figure 1.4), and thus, the DMFC 

suffers limitations from kinetically slow steps in these pathways. All simple 

alcohols considered for fuel cell use (CH3OH, CH3CH2OH, C2H2(OH)2, etc) react 

at low overpotentials to form aldehydes. For methanol, the H2CO generated is 

challenging to detect experimentally, but it is recognized as a significant by- 

product [1-4]. As a species formed in the early stages of methanol oxidation, 

H2CO accumulation lowers the energy conversion efficiency of the fuel. The 

primary in situ techniques used to study electrochemical reactions at the 

molecular level, infrared spectroscopy and mass spectroscopy, are not at all 

sensitive to H2CO. Therefore, in this project a sensitive fluorescence assay with 

good precision was adapted to detect and quantify H2CO. 

In the first part of the project, the cell that had been used in prior 

investigations of methanol oxidation pathways was improved by making it less 

permeable to atmospheric O2 and thereby less susceptible to interferences from 

O2 reduction in measurements at low electrode potentials. A MACOR glass 

119

ceramic disk was used in place of a Kel-F window that formed one side of a thin 

layer electrolysis cell. The MACOR was easily machined and enabled a fixed 

volume (50 µL) solution cavity to be created for the placement of sample during 

electrolysis. In contrast to Teflon or Kel-F, which are also easily machined and 

resistant to attack by most solvents, MACOR also possessed these properties 

and in addition did not allow the transport of atmospheric O2 into the cell. 

Methanol oxidation studies reported in this thesis were perfomed with the use of 

the MACOR cell. Final experiments reported demonstrate quantitative aspects of 

cell performance based on simple cyclic voltammetric measurements using 

hexaammine ruthenium (III) trichloride [Ru(NH3)6Cl3] as a redox probe. 

The cavity was shown to be somewhat thicker (~0.5 mm) than that for an Ideal 

thin layer cell. However, reversible, diffusion controlled cyclic voltammetry was 

achieved at scan rates of about 1 mV/s. 

The studies reported focused on applications of the cell in controlled 

potential electrolysis measurements of methanol oxidation over Pt and PtRu bulk 

metal and nanometer scale catalysts materials. Yields of H2CO produced during 

the electrolysis of 1.0 M methanol in 0.1 M H2SO4 for periods of 180 s were 

determined. In measurements with solid electrodes, it was observed that yields 

of H2CO greater than 10 % form at low potentials, especially over polycrystalline 

Pt, and can account for as much as 78% of the reaction charge. Figure 6.1 

compares the percent yield values graphically for H2CO produced during 

methanol electrochemical oxidation on all of the Pt and PtRu bulk electrodes 

studied. The most notable characteristic is the comparatively lower H2CO % 

120

yields (13% for XRu=0.3) for methanol oxidation on PtRu-solid compared to Pt 

solid electrode. Quantitative detection either for H2CO or CO2 becomes 

susceptible at lower electrode potentials on pure Pt catalysts, while on pure PtRu 

with optimum Ru composition (XRu=0.3) H2CO or CO2 can be assayed 

effortlessly at potentials below 0.6 VRHE. From these results, it can be concluded 

that the Ru metal with appropriate composition plays an important role in 

facilitating the complete oxidation of methanol. The dissociation of C-H bonds 

requires and ensemble of Pt sites. However, Ru sites lowers the potential 

(compared to Pt) for water activation and facilitates the conversion of methanol 

and any H2CO formed to more complete oxidation products, such as CO2 and 

formic acid (HCOOH) [5, 6]. It is predicted that Ru may break the H-OH bond in 

water more effectively than Pt to form the surface oxides necessary for the 

reaction to progress far along the path to CO2. In other words, Ru atoms 

dissociate water more easily than Pt to create oxygen species that lower H2CO 

production [6]. With higher concentration of Ru (XRu=0.9) in solid PtRu electrode 

the catalytic activity was inferior as it produces higher H2CO yield at all electrode 

potential compared to PtRu –solid (XRu=0.3) electrode. It is perhaps due to the 

surface oxides (RuOx) forming on Ru that are inactive for bond cleavage [5-8]. 

This may hinder the reaction and can cause large capacitance, which we 

observe in the cyclic voltammograms (Figure 3.1d, Figure 4.3a and Figure 4.3c) 

through out the “double layer charging region”. This is consistent with data 

previously reported [6, 8]. 

121

Following work with bulk electrodes, subsequent experiments extended 

the approach to investigate methanol electrochemical oxidation to H2CO on Pt 

and PtRu based nanometer scale catalyst materials. The results for all catalysts 

are summarized in Figure 6.2 graphically. Looking at Figure 6.1 and 6.2, it is 

noted that the H2CO yields with catalyst materials are considerably lower (below 

10% for the conditioned studied) than that of solid metal electrodes, which is 

consistent with earlier studies [9]. The CO2 measurements by IR spectrometry 

also established that the CO2 yield reached above 90% with rough catalyst 

surfaces [10]. This difference in H2CO yield on the bulk electrode and catalysts 

is believed to result from dissimilarities in the three dimensional structure of the 

electrode materials, and this topic will be discussed elaborately in the next 

section. It is also important to observe that the reaction charges for catalysts are 

significantly higher than those of bulk metal electrodes at all potentials (Table 3.1 

in Chapter III; Table 4.1 and 4.2 in Chapter IV). For example, at 0.8 VRHE, the 

highest reaction charge was 153 mC for nano particles whereas reactions over 

solid electrodes led to fairly low charges in the vicinity of 25 mC. The same 

pattern was observed for the lowest electrode potential (0.5 VRHE). The catalyst 

particles have nanometer size dimensions (~2-6 nm observed by transmission 

electron microscopy) with higher surface areas than the bulk electrodes. This 

higher surface area is at least partly responsible for the higher reaction charges 

recorded. The lower H2CO yields recorded for reactions over the catalyst 

particles likely result from the three dimensional nature of the catalyst bed, since 

partial oxidation products formed at particles deep within the bed can undergo 

122

Formaldehyde yield% 

90 

80 

70 

60 

50 

40 

30 

20 

10 

0 

Formaldehyde Yields for Oxidation of 1.0 M CH 3OH over Solid Bulk Electrode Materials 

Pt - solid 

PtRu - solid (XRu = 0.1) 



0.8 0.7 0.6 0.5 

E/V (vs. RHE) 

Figure 6.1: Bar charts comparing the formaldehyde yields from reaction of 1.0 M 

CH3OH in 0.1 M H2SO4 during 180 s electrolysis periods for solid bulk electrodes. 

123

Formaldehyde Yield% 

9 

8 

7 

6 

5 

4 

3 

2 

1 

0 

Formaldehyde Yields for Oxidation of 1.0 M CH 3OH over Nanoscale Catalyst Materials 

Pt-black 

C/Pt-10 wt% 

JM PtRu, XRu = 0.5 




SC PtRu, XRu = 0.1, 

with nafion 

0.8 0.7 0.6 0.5 

E/V vs. RHE 

Figure 6.2: Bar charts comparing the formaldehyde yields from reaction of 1.0 M 

CH3OH in 0.1 M H2SO4 over nanoscale catalyst materials during 180 s reaction 

periods. 

124

further reaction through interactions with other particles encountered along the 

diffusion path toward bulk solution. Furthermore, although the effect is small, the 

yields of H2CO appear to increase toward lower overpotentials at nanometer 

scale PtRu particles, similar to the trend displayed for the bulk electrodes. 

Comparing Figure 6.1 and Figure 6.2, there is a marked difference 

observed in the electrochemistry for methanol on bulk, solid electrodes and 

supported nanoparticle catalysts and it represents the more elaborate pictures of 

catalytic activity of the materials used for H2CO production. The spatial 

arrangement of particles and the diffusion conditions play critical roles in 

promoting methanol oxidation. The transport of reactants and products to and 

away from the electrode surface appears to be influenced by the particle layer 

properties as a concentration gradient is generated between the electrode 

surface and the reaction solution during electrolysis. For bulk metal electrodes, 

the transport of reactant occurs via planner diffusion, where reactant can only 

arrive following a perpendicular path relative to the electrode surface and can 

only react with the topmost atoms of the catalyst surface. After completing the 

reaction, the product can diffuse away to the bulk solution following the same 

perpendicular path and will most likely not encounter catalytic sites again. 

Chances that such an escape of the by-product without further oxidization will 

happen are much higher on smooth than on rough surfaces [11-14]. For 

nanoparticle metal cluster catalysts, there is a hemispherical diffusion where 

each site can behave similar to a small microelectrode, and the number of 

electroactive molecules having access to this hemispherical catalyst greatly 

125

exceeds that of a planner electrode [15]. From hemispherical volumes, the 

reaction analyte will be drawn around the electrode surface, and there is the 

possibility of more reactant-catalyst encounters as the reactant diffuses toward or 

away from the site [3]. The higher chances for multiple encounters move the 

methanol oxidation reaction to completion and can lead to lower yield of H2CO 

compared to solid metal electrodes. In other words, with multiple catalyst 

particles present in multilayer structure, readsorption of H2CO becomes 

increasingly more effective and increases the probability for the oxidation of 

volatile, oxidizable reaction intermediates to major products [2, 16-18]. There is 

another considerable hypothesis, which is the effect of the gradient diffusion 

layers close to and away from the particles surfaces. In the case of 

hemispherical diffusion, the gradient lines of the two particles diffusion patterns 

overlap. So there is a possibility that the reaction product spreads away from the 

first particle, it can encounter the other particle’s gradients and be pulled toward 

that site to react further. A more complete oxidation of the reactant should be 

occurred by this effect since byproducts again have more chances to encounter 

other catalyst sites [2, 3]. Actually this would cause to remove more H2CO from 

the solution and thus give the very low yields compared to that of solid 

electrodes. 

In the case of catalyst system, SC PtRu, XRu= 0.1 with Nafion, the nano 

catalyst structure is confined in a permeable film. The reaction environment is 

surrounded by three dimensional arrays of catalyst particles. As reaction 

proceeds, there is a spherical diffusion of reactants to each catalyst sites to 

126

eact. The partial products like H2CO, entrapping inside the three dimensional 

atmosphere, can just diffuse away from the site and feasibly meet another site for 

further oxidization to CO2. Thus the soluble partial oxidation products experience 

multiple encounters with the array of catalyst sites inside the film and the 

consequence is the increase of the chances for reactions to progress to 

completion with lower H2CO yield. 

The results obtained during this research project have opened several 

new avenues for research and development. The cell that has been used 

throughout these investigations can be further optimized with the modification of 

cell dimension which can offer reduced cell resistance and increased thin layer 

response. The proposed cell window with a novel MACOR material simplifies the 

cell fabrication process and the assembly of the spectrochemical cell. In the 

future, an integral piece of MACOR can be explored to construct the whole cell 

where the oxygen interference problem can be further minimized. The use of 

MACOR cavity can also be extended to quantify the other oxidation products like 

CO2 and HCOOH which can be detected off-line by IR spectrometry and 

chromatographic techniques. It will be very advantageous if the micro dimension 

cell can be utilized for studying the oxidation of CO with different electrode 

materials and the results can be compared with that of CO stripping process 

using conventional electrochemical cell. 

127

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4. Chen, Y. X.; Miki, A.; Ye, S.; Sakai, H.; Osawa, M. J. Am. Chem. Soc. 

2003, 125, 3680. 


Chem. 1993, 97, 12020. 

6. Rolison, D. R.; Hagans, P. L.; Swider, K. E.; Long, J. W. Langmuir 1999, 

15, 774. 

7. Long, J. W.; Rolison, D. R., In New Directions in Electroanalytical 

Chemistry II; Leddy, J., Vanysek, P. and Porter, M. D., Ed.; 

Electrochemical Society: Pennington, NJ, 1999; Vol. PV99-5, pp 125. 

8. Rolison, D. R.; Long, J. W.; Swider, K. E.; Merzbacher, C. I. Langmuir 

1999, 15, 780. 



11. Lu, G.-Q.; Chrzanowski, W.; Wieckowski, A. J. Phys. Chem. B 2000, 104, 

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12. Wang, H.; Loffler, T.; Baltruschat, H. J. Appl. Electrochem. 2001, 31, 759. 

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Electrochem. 1990, 20, 520. 

14. Lin, W.-F.; Wang, J.-T.; Savinell, R. F. J. Electrochem. Soc. 1997, 144, 

1917. 

15. Wightman, R.; Wipf, D. Electroanalytical Chemistry 1989, 15, 267. 


Electroanal. Chem. 2001, 509, 163. 

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17. Jusys, Z.; Kaiser, J.; Behm, R. J. Langmuir 2003, 19, 6759. 

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129

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