Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts [1 ed.] 9781617285639, 9781606927731

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Copyright © 2009. Nova Science Publishers, Incorporated. All rights reserved. Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

Copyright © 2009. Nova Science Publishers, Incorporated. All rights reserved. Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

POLYMER ELECTROLYTE MEMBRANE FUEL CELLS AND ELECTROCATALYSTS

Copyright © 2009. Nova Science Publishers, Incorporated. All rights reserved.

No part of this digital document may be reproduced, stored in a retrieval system or transmitted in any form or by any means. The publisher has taken reasonable care in the preparation of this digital document, but makes no expressed or implied warranty of any kind and assumes no responsibility for any errors or omissions. No liability is assumed for incidental or consequential damages in connection with or arising out of information contained herein. This digital document is sold with the clear understanding that the publisher is not engaged in rendering legal, medical or any other professional services.

Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

Copyright © 2009. Nova Science Publishers, Incorporated. All rights reserved. Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

POLYMER ELECTROLYTE MEMBRANE FUEL CELLS AND ELECTROCATALYSTS

RICHARD ESPOSITO AND

ANTONIO CONTI Copyright © 2009. Nova Science Publishers, Incorporated. All rights reserved.

EDITORS

Nova Science Publishers, Inc. New York

Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

Copyright © 2009 by Nova Science Publishers, Inc. All rights reserved. No part of this book may be reproduced, stored in a retrieval system or transmitted in any form or by any means: electronic, electrostatic, magnetic, tape, mechanical photocopying, recording or otherwise without the written permission of the Publisher. For permission to use material from this book please contact us: Telephone 631-231-7269; Fax 631-231-8175 Web Site: http://www.novapublishers.com NOTICE TO THE READER The Publisher has taken reasonable care in the preparation of this book, but makes no expressed or implied warranty of any kind and assumes no responsibility for any errors or omissions. No liability is assumed for incidental or consequential damages in connection with or arising out of information contained in this book. The Publisher shall not be liable for any special, consequential, or exemplary damages resulting, in whole or in part, from the readers’ use of, or reliance upon, this material. Independent verification should be sought for any data, advice or recommendations contained in this book. In addition, no responsibility is assumed by the publisher for any injury and/or damage to persons or property arising from any methods, products, instructions, ideas or otherwise contained in this publication.

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This publication is designed to provide accurate and authoritative information with regard to the subject matter covered herein. It is sold with the clear understanding that the Publisher is not engaged in rendering legal or any other professional services. If legal or any other expert assistance is required, the services of a competent person should be sought. FROM A DECLARATION OF PARTICIPANTS JOINTLY ADOPTED BY A COMMITTEE OF THE AMERICAN BAR ASSOCIATION AND A COMMITTEE OF PUBLISHERS. LIBRARY OF CONGRESS CATALOGING-IN-PUBLICATION DATA Polymer electrolyte membrane fuel cells and electrocatalysts / [edited by] Richard Esposito and Antonio Conti. p. cm. Includes index. ISBN 978-1-61728-563-9 (E-Book) 1. Proton exchange membrane fuel cells. 2. Electrocatalysis. I. Esposito, Richard, 1963- II. Conti, Antonio, 1962TK2931.P65 2009 621.31'2429--dc22 2009006861

Published by Nova Science Publishers, Inc.    New York

Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

CONTENTS Preface Chapter 1

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Chapter 2

vii Synthesis of Polymer Electrolyte Membrane for Fuel Cell Applications Susanta K. Das

1

The Development of Bipolar Plate Materials for Polymer Electrolyte Membrane Fuel Cells (PEMFC) Yan Wang and Derek O. Northwood

33

Chapter 3

Fuel Cell Converters for High Power Applications Phatiphat Thounthong, and Bernard Davat

Chapter 4

CFD Models for Analysis and Design of Ambient Air-Breathing PEM Fuel Cells Maher A.R. Sadiq Al-Baghdadi

Chapter 5

New Trends in the Development of PEMFC Catalysts Alevtina L. Smirnova

Chapter 6

Air-Breathing Direct Methanol Fuel Cells with Catalysed Titanium Mesh Electrodes Raghuram Chetty and Keith Scott

Chapter 7

Metal Electrocatalysts for Direct Liquid-Feed Fuel Cells Umit B. Demirci

Chapter 8

The Electro-Catalytic and Mass Transport Components of the Electrode Potential Loss in a PAFC Cathode Determined Using a Numerical Model G. Psofogiannakis,, Y. Bourgault, B.E. Conway and M. Ternan

Chapter 9

Alcohol Oxidation on Pt and Pd Based Electrocatalysts Changwei Xu and Yexiang Tong

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63

123 167

193 205

239

275

vi Chapter 10

Chapter 11

Chapter 12

Contents Physicochemical and Electrocatalytic Properties of PtRu/C Prepared by Impregnation Reduction Method: Effect of Preparation Parameters Jun Li, Jing-Dong Lin and Dai-Wei Liao

293

Conducting Polymers Used as Catalyst Support for Fuel Cell Application Lei Li and Jun Yang

309

Electrocatalytic Modified Electrodes with Transition Metal Azamacrocycles and Other Complexes for the Detection of Sulfur and Nitrogen Oxoanions Maria J. Aguirre, Yo-ying Chen, Galo Ramirez, Mauricio Isaacs, Fabiola Isaacs and William Cheuquepan

341

421

Index

439

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Short Communication: Compound Micro-Grid of Hydrogenation City-Gas Engine and PEM Fuel Cell Shin’ya Obara

Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

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PREFACE This book presents current research in fuel cells which are growing in importance as sources of sustainable energy and are forming part of the changing program of energy resources. Fuel cells provide environmentally friendly, clean and highly efficient energy source for power generation. In order to efficiently utilize the energy from fuel cells, a power conditioning system is required. This book describes the converters’ basic operating principles and analyzes performance for low-voltage, high-power fuel cell applications. Full three-dimensional, multi-phase, non-isothermal computational fluid dynamics models of planar and novel tubular-shaped air-breathing proton exchange membrane fuel cell are also presented in detail. Research and review of electrocatalysts such as platinum are presented as well. Chapter 1 - In this chapter, we discuss state of the art approach to the design and fabrication of proton exchange membrane (PEM) and the approach we have developed whereby a non-structural polymer fabricated for proton exchange capacity is bound to an inert polymer matrix. Our fabrication technique separates the proton exchange and structural requirements of the proton exchange membrane allowing greater flexibility in proton exchange membrane design. Results related to performance of novel PEMs are presented here and thus omitting structural characterizations due to limited chapter space. The proton exchange material described herein is a ter-polymer composed of various ratios of acrylic acid, styrene and vinylsulfonic acid. These materials were bound to an inert ethylenetetrafluoroethylene copolymer mesh that had been rendered adhesive by using patented hydroxylation techniques in a two-step water-borne process. The basic characteristics of our new membrane, related only to the performance of proton exchange membrane for fuel cell applications, were compared to those of Nafion® 212 and benchmarked the performance of the membranes. Results indicated that besides improvement of all physical quantities, the new membrane could transfer protons approximately 10 times faster than Nafion® 212 per unit area under the test conditions utilized at 80oC. Reduction of membrane’s relative resistance is also achieved. Besides low induction time, the in-house membranes are able to conduct protons at reduced membrane water content at 80oC compared to Nafion® 212. Chapter 2 - With escalating oil prices and heightened environmental concerns, increasing attention is being paid to fuel cell technology. Polymer Electrolyte Membrane Fuel Cells (PEMFCs) have significant advantages compared with other types of fuel cell. PEMFCs have a solid electrolyte which provides excellent resistance to gas crossover. PEMFCs are considered as promising candidates for transportation applications because of their lower

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viii

Richard Esposito and Antonio Conti

temperature operation and fast start-up. However, there are a number of barriers, e.g. price, weight and volume, which have prevented PEMFCs from being more widely used. Bipolar plates (BPs) are very important components in the PEMFC stack. They are designed to accomplish many functions, including: distribute the fuel and oxidant in the stack; facilitate water management within the cell; separate the individual cells in the stack; carry current away from the cell; and facilitate heat management. Currently, BPs are made of non-porous graphite because of its chemical and thermal stability in the PEMFC working conditions. However, their high price and the difficulties in machining the gas flow channels present problems for the application of non-porous graphite BPs. Also, non-porous graphite is very brittle, and is therefore readily broken, with the result that H2 and O2 could leak into the whole fuel cell system, which is potentially dangerous for fuel cell operation. Increasing attention is, therefore, being paid to the use of metallic and composite materials as replacements for non-porous graphite in BPs for PEMFCs. The ideal BP material should demonstrate high values of electrical conductivity, thermal conductivity, corrosion resistance and compressive strength and low values of gas permeability and density. In this paper, the three types of BP materials, i.e. non-porous graphite, metallic and composite materials, are reviewed for their physical, chemical and electrochemical properties. For the metallic BPs, a wide variety of coating techniques and materials have been used in order to increase the corrosion resistance of the base materials. For the composite materials, different composite material compositions and processing methods have been used to increase the electrical conductivity and production rates. The most promising candidate materials for BPs that meet the design requirements are identified. Chapter 3 - Fuel cells (FCs) provide a clean and highly efficient energy source for power generation; however, in order to efficiently utilize the energy from fuel cells, a power conditioning system is required. Typical fuel cell systems for stand-alone and utility grid-tied stationary power applications are found mostly with low nominal output voltages around 24 V, 48 V, 400 V or 750 V and power levels are found to be 0.5 kW to 2 MW. A power conditioning system for such applications generally consists of a dc/dc converter and a dc/ac inverter, and the dc/dc converter for low-voltage, high-power fuel cells must deal with a high voltage step-up conversion ratio and high input currents. Although many dc/dc converters have been proposed, most deal with high input voltage systems that focus on step-down applications, and such dc/dc converters are not suitable for low-voltage, high-power fuel cell applications. With the development of fuel cell based power systems, the need for more advanced dc/dc power converters has become apparent. In such applications dc/dc converters provide an important link between low voltage fuel cell sources and inverter buses operating at significantly higher voltages. Advancements in converter efficiency, cost reduction, and size reduction are the most necessary. These challenges are formidable, even when considering the improvements made to conventional dc/dc topologies. However, it can be possible to achieve these criteria through the implementation of more advanced topologies. Single phase or multiphase non-isolated or isolated dc/dc converters have been studied in this paper. For multiphase power converters, they offer several advantages that are very desirable in low-voltage, high-power fuel cell applications. First, a multiphase is constructed with paralleled phases, which increase power rating and current handling capability for high input current. Second, an interleaving control scheme produces a high operating frequency

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Preface

ix

with a low switching frequency, and the high operating frequency reduces size of passive components. A recently developed efficient multiphase dc/dc topology offers benefits over standard designs. Passive component sizes and output ripple voltage were reduced as a result of an effective boost in switching frequency. For cost reduction, the converter was designed and built with discrete components instead of more expensive integrated modules. This chapter examines several dc/dc converter topologies that are targeted for fuel cell applications. The paper also describes the converters’ basic operating principles and analyzes performance for low-voltage, high-power fuel cell applications. 1.2-kW prototypes for each converter are built and tested with a real fuel cell system. Moreover, the characteristics of three fuel cell systems of 0.5-kW, 40-A PEM fuel cell; 0.5-kW, 50-A PEM fuel cell; and 1.2kW, 46-A PEM fuel cell in our laboratories in Nancy-Lorraine, France and in Bangkok, Thailand have been presented when the fuel cells connect with power converters at a high switching frequency. Chapter 4 - Fuel cell system is an advanced power system for the future that is sustainable, clean and environmental friendly. Fuel cells are growing in importance as sources of sustainable energy and will doubtless form part of the changing programme of energy resources in the future. Small fuel cells have provided significant advantages in portable electronic applications over conventional battery systems. Competitive costs, instant recharge, and high energy density make fuel cells ideal for supplanting batteries in portable electronic devices. For portable applications like laptops, camcorders, and mobile phones the requirements of the fuel cell systems are even more specific than for vehicle applications. The requirements for portable applications are mostly focused on size and weight of the system as well as the temperature. Among all kinds of fuel cells, proton exchange membrane (PEM) fuel cells are compact and lightweight, work at low temperatures with a high output power density and low environmental impact, and offer superior system start-up and shutdown performance. However, the typical PEM fuel cell system with its heavy reliance on subsystems for cooling, humidification and air supply would not be practical in small applications. The air-breathing PEM fuel cells without moving parts (external humidification instrument, fans or pumps) are one of the most competitive candidates for future portablepower applications. The development of physically representative models that allow reliable simulation of the processes under realistic conditions is essential to the development and optimization of fuel cells, the introduction of cheaper materials and fabrication techniques, and the design and development of novel architectures. The difficult experimental environment of fuel cell systems has stimulated efforts to develop models that could simulate and predict multidimensional coupled transport of reactants, heat and charged species using computational fluid dynamic (CFD) methods. The strength of the CFD numerical approach is in providing detailed insight into the various transport mechanisms and their interaction, and in the possibility of performing parameters sensitivity analyses. The results of CFD analyses are relevant in: conceptual studies of new designs, detailed product development, troubleshooting, and redesign. CFD analysis complements testing and experimentation, by reduces the total effort required in the experiment design and data acquisition. In this Chapter, full three-dimensional, multi-phase, non-isothermal computational fluid dynamics models of planar and novel tubular-shaped air-breathing PEM fuel cell have been developed and presented in detail. These comprehensive models account for all major

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1940s [104-106]. It is an energy carrier when it directly powers a PEMFC-type fuel cell, socalled direct borohydride fuel cell (DBFC) [103, 107]; it is a hydrogen carrier when it stores atomic hydrogen (and generates hydrogen to feed a PEMFC) [108]. In other words, NaBH4 is able to feed directly or indirectly a PEMFC (Figure 9). The grey square should include > In the pink area, it is >

Aqueous alkaline solution of NaBH4 Oxidation

Hydrolysis Direct fuel

Indirect

H

Fuel cells

Direct borohydride fuel

H2-powered PEMFC

Figure 9. Borohydride oxidation (blue area) versus borohydride hydrolysis (pink area) for the PEMtype fuel cells.

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In the 1940s, sodium borohydride attracted attention because of its ability to generate molecular hydrogen, the generation reaction being spontaneous and besides being able to be accelerated thanks to either acids or metal-based catalysts like e.g. cobalt or ruthenium [105, 109, 110]. In the meantime, the NaBH4 oxidation began drawing some researchers’ attention [111]. In 1953, Pecsok [112] proposed for the polarographic oxidation of NaBH4 the net reaction (9) and calculated the standard electromotive force of this half reaction as being 1.23 V versus the standard hydrogen electrode: BH4- + 8 OH- Æ BO2- + 6 H2O + 8 e-

(9)

In 1962, Indig and Snider [113] proposed the concept of the DBFC because the possibility of using NaBH4 as an anodic fuel appeared attractive due to its extremely low equivalent weight of 4.73 g. However it was concluded that the possibility of achieving a reversible NaBH4 electrode with an 8-electron oxidation appeared unlikely and that one might have to settle for a reaction involving 4-electron change (10) and a lower potential: BH4- + 4 OH- Æ BO2- + 2 H2 + 2 H2O + 4 e-

(10)

From the 1960s to the late 1990s, NaBH4 was not considered in the energy area anymore. However, since 1999, the year of Millenium Cell’s articles publishing, NaBH4 attracted attention one more time and today increasingly interests as energy/hydrogen carrier. DBFC is based on the borohydride oxidation. The oxidation reaction takes place under an alkaline medium because the borohydride ions are not chemically stable in acidic media. Basically, these ions self-hydrolyze generating molecular hydrogen according to the theoretical reaction (11) that follows:

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Preface

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considerable hydrodynamic and mass transport limitations. To circumvent these problems, an expanded titanium mesh has been adopted as the catalyst substrate material in this study. Titanium mesh was used as the substrate due to its chemical stability and its ability to support a diverse range of electrocatalysts. In the proposed fuel cell application, the mesh-based electrode has several potential advantages in terms of cost, simplicity, size and shape. This work describes the design, fabrication and evaluation of a passive air-breathing direct methanol fuel cell using the mesh-based electrodes. PtRu/Ti and Pt/Ti prepared by electrodeposition onto the Ti mesh were used as anode and cathode, respectively. Methanol is stored in an in-built reservoir and oxygen is taken from the surrounding air. Single cells with an active area of 9 cm2 produced a power density of 9.5 mW cm-2, whereas a cubic four-cell stack with an active area of 24 cm2 produced a maximum power of 180 mW. The effects of experimental parameters such as concentration and temperature on the cell performance were also investigated. Chapter 7 - Direct liquid-feed fuel cells (DLFCs) that are intended for mobile/portable devices are technologies based on the polymer electrolyte membrane fuel cell (PEMFC). The technical, commercial development of DLFCs is dependent on various materials like, e.g., metal electrocatalysts. These electrocatalysts are key materials. They catalyze the fuel oxidation as well as the oxygen reduction. The main difference between DLFCs and PEMFC is the fuel, that is, liquid chemicals and hydrogen gas respectively. The early investigations about DLFCs assessed Pt electrocatalysts that were being used for the PEMFC but with the fuels being different, new issues arose. For instance, in the direct methanol fuel cell (DMFC), Pt poisoned because of the formation of CO-like intermediates like CO, formaldehyde or formic acid. Hence new metal-based catalytic materials were regarded and are today being investigated. This is discussed in the present chapter. Many liquid chemicals have been envisaged as fuels of DLFCs but the paper only deals with the main ones: i.e., the DLFCs fed with, e.g., C1-C3 alcohols, formic acid and sodium borohydride. In the DLFCs, whatever the fuel (e.g., methanol, ethanol, formic acid, sodium borohydride...) is, the anode electrocatalysts are generally metal-based. Monometallic electrocatalysts have been widely investigated and, today, lots of new bimetallic (and even trimetallic) combinations are suggested throughout the open literature. Globally, the investigations devoted to the DLFCs electrocatalysts focus on reactivity and durability. With the multimetallic electrocatalysts, improved reactivity and durability may be obtained because of electronic and/or geometric effects. In other words, the addition of a metal (e.g., Ru) to another (e.g., Pt) may modify the electronic and geometric features of the latter, influencing then its adsorption properties, reactivity and durability. This, especially, is discussed in the present chapter. Theoretical results/data (e.g., d-band centre of metals and segregation, both from density functional theory) are useful tools for trying to understand the catalytic behaviors of metal electrocatalysts. These tools, available in the open literature, are used in the present discussion for interpreting electronic and geometric effects. Furthermore, when possible these are exploited for suggesting alternative multimetallic combinations as potential anode electrocatalysts for the DLFCs that are tackled. Chapter 8 - The electro-catalytic and mass transport components of the electrode potential loss in the cathode of a phosphoric acid fuel cell (PAFC) were investigated through the development a numerical model. A one-dimensional model was used to describe the porous electrode that was composed of agglomerates containing platinum catalyst / carbon /

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polytetrafluoroethylene components. The porous structure of the spherical agglomerates was partially filled with phosphoric acid electrolyte while a liquid film of electrolyte surrounded the agglomerates. The system of equations that described the system was solved using the shooting method. The predictions made by the model were in good agreement with two different sets of experimental polarization data. Increasing the loading of the platinum electro-catalyst from 0.01 to 1 mg/cm2 increased the current density by one and one-half orders of magnitude. Increasing the thickness of the electro-catalyst layer from 50 μm to 250 μm increased the current density by a factor of three. Decreasing the radii of the electrocatalyst containing agglomerates from 5 μm to 0.1 μm also increased the current density by a factor of three. The oxygen reactant concentration profiles, from the agglomerate exterior to its center, were determined. The oxygen concentration at the center was essentially zero in agglomerates having a diameter of 2 μm, whereas it was 73% of the exterior concentration in agglomerates having a diameter of 0.5 μm. In addition, the model identified the fraction of the electrical potential loss caused by each of the following phenomena; electro-catalytic surface processes (cathode overpotential), ionic conduction, diffusion of oxygen gas, and diffusion of oxygen dissolved in the liquid electrolyte. Of these, the electro-catalytic surface processes contributed the largest single electrical potential loss. Chapter 9 - Pt has been extensively investigated as the elctrocatalyst for methanol and ethanol electrooxidation in acid media. Some attention has paid to 1-propanol, 2-propanol and EG electrooxidation in acid media. The methanol gives the best performance and 2-propanol has very high current at low potential on Pt in acid media. When the alcohol electrooxidation operates in alkaline instead of acidic media the activity for all alcohols oxidation is significantly improved. EG shows the best electrooxidation activity on Pt in alkaline media. Pd is not a good electrocatalyst for methanol and EG, but it shows excellently higher activity than Pt for 1-propanol, 2-propanol and ethanol electrooxidation in alkaline media. For decades, significant progress has been made in the development of other metal modified Pt and Pd electrocatalysts for alcohol oxidation. It is well known that the Pt alloys are Pt-Ru and Pt-Sn and the correlated ternary Pt-Ru-based and Pt-Sn-based catalysts. The PtNi and PdNi have attracted more attention for alcohol electrooxidation. Enhanced activity of alcohol electrooxidation such as lower onset potential and improved stability was responsible for the change in the electronic properties of Pt and Pd in PtNi and PdNi. We have dispersed the PtNi and PdNi nanoparticles on Ti foil with a porous structure which were successfully prepared by electrodeposition method. Aaddition of Ni significantly promotes catalytic activity of the Pt and Pd electrocatalysts for the methanol and ethanol electrooxidation. Oxide promoted Pt and Pd catalysts for methanol and ethanol electrooxidation have been extensively investigated. In the oxides, the CeO2 and WO3 have attracted enormous attention. We have studied the alcohol electrooxidation on Pt-oxide and Pd-oxide in alkaline media. The addition of oxide to Pd promotes significantly the catalytic activity for the ethanol oxidation. The onset potential shifts the negative direction and the current density increases for the ethanol oxidation reaction on Pd-oxide/C in comparison to that on Pd/C. Chapter 10 - The impregnation reduction method was used to investigate the influence of preparation parameters on the physicochemical and electrocatalytic properties of 40 wt% PtRu/C. Based on the experimental results, it is found that the reductants with higher reducing power, such as formaldehyde, KBH4, methanol, N2H4, lead to smaller particles and higher

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Preface

xiii

dispersion. On the other hand, the reductants with lower reducing power, such as ethanol, isopropyl alcohol, result in bigger particles and poorer dispersion. In addition, the effect of pH value during impregnation was also investigated. In the case of using NaBH4 as a reductant, the favorable pH value is below 4, whereas in the case of using methanol, N2H4 and formaldehyde as the reductant, the favorable pH value is higher than 8. The characterization of the 40 wt% PtRu/C catalysts with X-ray diffraction (XRD) and transmission electron microscope show that the catalysts prepared under the condition of proper pH and α values (stoichiometric ratio of reductant to Pt and Ru metals) have welldispersed PtRu particles with average diameter around 2.3 nm. XRD experiments also indicate that the PtRu/C catalyst prepared by the reductant with different stoichiometric ratios and different pH values has different degrees of PtRu alloying. The effects of pH of precursor solution and α values of reductant on the average particle size of PtRu/C are discussed in detail. The electrocatalytic activity of the as-prepared catalysts was tested by cyclic voltammograms both in 1 M H2SO4 and 0.5 M H2SO4+ 1 M CH3OH. The results indicate that a reductant with higher reducing power, a high stoichiometric ratio of the reductant and a proper pH value are favorable to synthesize a catalyst with a higher electrocatalytic activity toward methanol oxidation. In addition, a new and highly reproducible method based on the flow injection (FI) technology was also developed to synthesize PtRu/C automatically. This method was proven to be a good fit for the mass production of PtRu/C. These benefits include effective, raw material saving, highly reproducible and labor-saving. Chapter 11 - An ideal catalyst support for fuel cell application has not only high electron and proton conductivity, but also good permeability to gases and water. However, carbon, the most common catalyst support used in fuel cell, is impermeable to gases (oxygen, hydrogen and water vapor) and non-conductive to protons, which results in low catalyst utilization and poor catalyst performance. Recently, more attention has focused on the use of conducting polymers (e.g. polyaniline, polypyrrole and polythiophene) as catalyst supports and the promising results have been obtained. Combined with our recent research results, a review on conducting polymers and conducting polymers-carbon composites used as catalyst support for fuel cell application will be presented in this chapter. Chapter 12 - Modified electrodes with transition metal phenantrolines, phthalocyanines, porphyrins and other complexes where the metal center is a transition metal have been widely studied because they show highly activity toward the oxidation and reduction of several analytes. One of the most interesting applications of these modified electrodes is the determination of oxoanions in waste water, food and beverages. In fact, they have been used to detect nitrite, nitrate, sulfite, and other pollutants or traces with low limit detection and in some cases; specificity. On the other hand, it is interesting that these modified electrodes can be designed for desirable purposes by changing the metal center, modifying the ligand, or changing the electrolyte. There are some cases in which the position of a substituent on the ligand can drastically change its activity. The modified electrodes are been used as supramolecular assemblies, as electroactive films or as composites of different layers adsorbed on the surface. This brief review shows the last decade’s studies on the detection of sulfur and nitrogen oxoanions using electrodes modified with metal transition complexes. Short Communication - The independent micro-grid is considered to be a technology in which maximum distributed energy is realizable. However, there are many subjects, such as the stability of the dynamic characteristics of power and development of an optimal design method. If the fuel cell system of the capacity corresponding to a load peak is installed,

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equipment cost will be high and energy cost will not be able to get any profile commercially. By increasing the hydrogen concentration at the time of low load, the power-generation efficiency of a city-gas-engine-generator improves, and carbon dioxide emissions decrease. So, in this Chapter, a micro-grid composed from a proton exchange membrane fuel cell and a hydrogenation city gas engine was investigated using numerical simulation. The system with a small load factor of NEG and with a large load factor of PEM-FC system has few CO2 emissions. The system which combined base-load operation of PEM fuel cell and load fluctuation operation of hydrogenation city gas engine is the most advantageous for the comprehensive evaluation of equipment cost, power generation efficiency, and CO2 emissions. When the optimal system was installed into the urban area model of 20 buildings and analyzed, power generation efficiency was 25% and CO2 emissions were 1,106 kg/Day.

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In: Polymer Electrolyte Membrane Fuel Cells… Editors: R. Esposito, A. Conti

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Chapter 1

SYNTHESIS OF POLYMER ELECTROLYTE MEMBRANE FOR FUEL CELL APPLICATIONS Susanta K. Das* Department of Mechanical Engineering Center for Fuel Cell Systems and Powertrain Integrations Kettering University, Flint, Michigan-USA

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ABSTRACT In this chapter, we discuss state of the art approach to the design and fabrication of proton exchange membrane (PEM) and the approach we have developed whereby a nonstructural polymer fabricated for proton exchange capacity is bound to an inert polymer matrix. Our fabrication technique separates the proton exchange and structural requirements of the proton exchange membrane allowing greater flexibility in proton exchange membrane design. Results related to performance of novel PEMs are presented here and thus omitting structural characterizations due to limited chapter space. The proton exchange material described herein is a ter-polymer composed of various ratios of acrylic acid, styrene and vinylsulfonic acid. These materials were bound to an inert ethylene-tetrafluoroethylene copolymer mesh that had been rendered adhesive by using patented hydroxylation techniques in a two-step water-borne process. The basic characteristics of our new membrane, related only to the performance of proton exchange membrane for fuel cell applications, were compared to those of Nafion® 212 and benchmarked the performance of the membranes. Results indicated that besides improvement of all physical quantities, the new membrane could transfer protons approximately 10 times faster than Nafion® 212 per unit area under the test conditions utilized at 80oC. Reduction of membrane’s relative resistance is also achieved. Besides low induction time, the in-house membranes are able to conduct protons at reduced membrane water content at 80oC compared to Nafion® 212.

*1700 West University Avenue, Flint, MI 48504 Tel.: (810)-762-9916 FAX: (810)-762-7860 Email: [email protected]

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Susanta K. Das

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1. INTRODUCTION Fossil fuel combustion has contributed significantly to the accumulation of atmospheric greenhouse gases for environment pollution causing health problems in many urban areas and also contributed significantly in detrimental global climate changes. For the sake of sustainable development and environmental protection, it is essential to reduce the alarming consumption rate of finite fossil fuel reserves and search for alternative fuels and energy sources. Having almost no emission except for water and heat as by-product, fuel cell (FC) an electrochemical device, is a promising candidate for a potential alternative, renewable, clean and energy efficient power source. Proton exchange membrane fuel cell (PEMFC) is a good candidate among different types of fuel cells [1] for providing an alternative and sustainable clean energy for remote power supplies, potable power devices, stationary power generations, automotive power systems and a wide range of transportation applications [1-3]. In PEMFCs, a membrane (electrolyte) is required to separate the chemical reactions at the electro-catalytic porous anode and cathode electrodes [1]. The protons produced in chemical reactions at anode travel through the membrane by diffusion/migration to the cathode for the electro-chemical reactions to take place. These membranes have fixed anionic charges which often permit easy proton transport. The commercial success of PEMFCs depends more on the cost effectiveness of its components such as membrane, catalyst, bipolar plate materials, etc as compared to other energy conversion and power generation devices. Cost reduction of PEM fuel cells can be achieved by enhancing performance and lowering material and fabrication costs. Currently a benchmark commercially produced and widely used proton exchange membrane (PEM) for fuel cell applications is Nafion® [4-11]. Nafion® is a sulfonated fluoropolymer with the chemical structure indicated in Figure 1. Nafion has a number of limitations such as an operating temperature range of 50oC -90oC [5-11], undesirable reactant permeability - on the order of 10-6 cm2/s [9-14] which results in decreased fuel cell performance, limited operational hydration range [1-4, 11-13] and high cost, US $800 per m2 [1-4]. To overcome these limitations, a variety of alternative materials for PEM fabrication have been explored.

CF2 CF2

CF2

x

CF

y n CF2

CF2 CF CF3

O- H+

CF2

z

CF2

S O

Figure 1. Skeleton of chemical composition of Nafion®.

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For example, “self humidifying” poly-tetrafluoroethylenes (PTFEs) polymer have been developed [5-11]. Furthermore, strong acids such as phosphoric acid, sulfuric acid, and alphaF sulfuric acids [5-14] have been added to a variety of materials to improve proton conductivity. Among these, the most common proton transfer functional group has been the sulfonate group [4-6]. The proton conductivity of acid-doped membranes can be several orders of magnitude greater than that of Nafion® but unbound acid species are rapidly leached from the membrane in tens to hundreds of hours. The essential requirements of PEMs for fuel cell applications include: (i) high proton conductivity, (ii) minimal thickness (to minimize resistance resulting in fuel cell’s ohmic drop), (iii) high thermal stability, (iv) excellent mechanical properties (strength, flexibility, and processability), (v) excellent chemical stability, (vi) low water drag, (vii) rapid adjustment of fast kinetics for electrode reactions, (viii) low or minimal gas permeability and finally (ix) low cost and high availability. A common polymer backbone used in fuel cell membrane preparation is polystyrene (PS) [6-8]. Chemical structure of sulfonated PS is exhibited in Figure 2. This polymer has been doped with sulfonate functionality [6-7] or been sulfonated directly [4-7]. Direct sulfonation greatly increases the stability of the sulfonate groups within the PEM and thus greatly increases membrane life expectancy. Hence, this type of membrane has contributed greatly to the use of sulfonated polymers as fuel cell electrolytes [7-9]. These sulfonated PS based PEMs have disadvantages of their own including lower C-H bond (chemical) stability compared to perfluoronated PEMs [4-9]. Sulfonated PS based polymer membranes further suffer from a short lifetime relative to their perfluoronated PEM counterparts. This is due to the sensitivity to oxidation of the tertiary C-H bonds in the styrene chains by oxygen and hydrogen peroxide under fuel cell operational conditions. Chemically, the bond strength for C-F is about 485 kJ per mole, much higher than that of C-H bonds (typically 350-435 kJ per mole) or C-C bonds (typically 350-410 kJ per mole) [13].

O O

O- H+ S

n Figure 2. General chemical structure of sulfonated polystyrene.

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Susanta K. Das

Polymers containing C-F bonds have also been shown to have high thermal and chemical stability [1-4, 13-15]. For example, PTFEs, consisting of repeated unit of –[CF2-CF2]-, have an excellent chemical stability and a high melting temperature around 370oC relative to the sulfonated PS based PEMs [13]. In attempts to prepare a more structurally robust PEM, additives such as silica have been used [5-8]. The addition of other additives has sometimes been for the purpose of providing an anchor for the PEM [6-8]. Ethylene-tetrafluoroethylene (ETFE) film has similar properties to the current fluorinated polymers used to prepare PEMs and thus has been used to provide a robust matrix from which sulfonate based PEMs were prepared. These copolymer films are currently prepared through radiation grafting [14-15]. For example, toluene and styrene have been grafted onto ETFE film, with and without divinylbenzene cross-linker [15]. These grafted films were then sulfonated using chlorosulfonic acid to provide proton exchange capability [13-15]. Although in many aspects, these membranes performed at least as good as Nafion®, their manufacture involves the use of both radiation and chlorosulfonic acid and thus will not likely be industrially feasible on a large scale. Although perfluorosulfate based membranes have steadily advanced in terms of fuel cell's performance, one major drawback of these membranes is their low conductivity and thus poor cell performance under low humidity and elevated temperatures (above 90oC) due to water loss [6-7]. Many additional polymer types have been explored for the preparation of PEMs. For example, polysiloxane polymers have been examined [16-19]. Furthermore aromatic polymers containing a phenylene backbone have been explored [20]. Other fluorinated polymer membranes have also been actively investigated, for example, polytetrafluoroethylene-hexafluoropropylene (PTFEP) films [21-23]. The PTFEP film is irradiated first, and then styrene groups are grafted on it with divinylbenzene (DVB) as a cross-linker. In this material, the proton conductivity is introduced by sulfonating the aryl groups. A recent work using this type of membrane has reported a fuel cell lifetime over 5000 hours at 85oC [23]. Poly-vinylidene fluoride (PVDF) based polymer membrane has also been developed through grafting and then sulfonating the styrene groups [24-27]. In these PVDF based PEMs, the PVDF provides good physical stability and chemical resistance while the sulfonated polystyrene provides good conductive properties and high water uptake [28-30]. However, high conductivity is only obtained at high levels of sulfonation which also results in high swelling and thus poor mechanical properties, especially at higher operating temperatures [31]. Great efforts have been made to reduce the swelling by ionic cross-linking [32-35]. Another method used to maintain high conductivity while preventing swelling has been to prepare inorganic-organic composites [36-40]. Inorganic-organic composite membranes have been prepared by (i) casting a bulk mixture of powder or colloidal inorganic components with a polymer solution, or (ii) in-situ formation of inorganic components in a polymer membrane or in a polymer solution. The bulk polymer solution enables the nanosized particles to be dispersed in the formed membranes [41]. Some of these composite membranes exhibit promising conductivities at temperatures above 100oC [20, 42]. However, most of these composite membranes have not been tested in fuel cell applications. Acid-base polymers are another class of proton exchange membranes which provide good performance at high temperatures [20, 43]. One of these, phosphoric acid-doped poly[2,2’-(m-phenylene)-5,5’-benzimidazole] (PBI), ionically cross-linked has recently received much attention [33, 44]. In these acid-base polymer membranes, basic polymers can be doped with an amphoteric acid, which acts both as a proton donor and acceptor and

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therefore allows for proton migration. It was found that the base protonation and strong hydrogen bridging in acid-base blend membranes remarkably reduced polymer swelling [32]. The resulting acid-base blends constitute a new class of proton conducting membranes with high conductivity, thermal stability, and mechanical flexibility and strengths [44]. In general, polymers bearing basic sites such as ether, alcohol, imine, amide or imide groups react with strong acids such as phosphoric acid or sulfuric acid. The basic components of these polymers enable the establishment of hydrogen bonds with the acid. Because of their unique proton conduction mechanism by self-ionization and self-dehydration, H3PO4 and H2SO4 exhibit effective proton conductivity even in an anhydrous form. When basic components are present, the interaction between basic sites and the acids increase extent of the acid dissociation relative to that of the anhydrous acids [45]. Most of the studied acid/polymer systems are not entirely anhydrous, as water is present as a necessary plasticizer for improving conductivity and mechanical properties [45]. The acid-base polymeric electrolyte has been proposed for fuel cell membrane electrolyte at temperature above 100oC [46-47]. Acid-base membrane has been cast from solutions of different concentrations of organic solvents, for example, 2.5-3.0% solution in a mixture of NaOH and ethanol [48]. The cast membrane is then doped with the acid in order to obtain sufficient conductivity [45]. Sulfonated polysulfone/PBI (SPSF/PBI) membranes doped with phosphoric acid have been investigated and shown to exhibit excellent chemical and thermal stability as well as good proton conductivity [49]. Using these materials, good fuel cell performance has been achieved at a temperature of 110oC and pressure of 1.5 atm [50-51]. The objective of the present study is to develop a novel, efficient and inexpensive proton exchange membrane (PEM) for fuel cell applications that is comprised of a proton exchange media cast onto a robust polymer mesh to form the hybrid PEM. In this manner, many of the membranes stability requirements can be met with the polymer mesh support while the proton exchange media may be tailored to most efficiently exchange protons. The polymer mesh used here is a surface modified ethylene-tetrafluoroethylene (ETFE) copolymer, which has similar properties as of fluorinated polymers [1-8], In this study, a novel water-borne chemical procedure was used to functionalize ETFE mesh onto which proton conducting polymer was casted with varying compositions of random ter-polymers comprised of acrylic acid, styrene and vinylsulfonic acid. To evaluate the performance of these hybrid PEMs, an easy and inexpensive test method was developed. Details of materials and methods are discussed in section 2. Experimental details for performance evaluation of the novel membranes are presented in section 3. Results and discussions are provided in section 4 and finally conclusions, based on the results obtained for the new membrane compared to Nafion® 212 membrane, are drawn in section 5.

2. MATERIALS AND METHODS Many types of polymer materials and their synthesis methods have been explored in recent years for the preparation of PEMs [1-32]. Ultimate aim of synthesis of polymer materials is to get the best performance of proton exchange membrane materials in operational fuel cell environment. The materials and manufacturing procedures used to develop our in-house proton exchange membrane for fuel cell applications are described below.

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Susanta K. Das

2.1 Materials Ethylene-tetrafluoroethylene (ETFE) mesh (70µm nominal aperture, 66.7 threads/square inch, 70 µm monofilament diameter, 21% open area) was purchased from Goodfellow Corporation, Oakdale, PA, U.S.A.. 15% sodium hypochlorite and 85% phosphoric acid were purchased from PVS Nolwood, Detroit, MI, U.S. A.. All other chemicals and polymers were purchased from Sigma-Aldrich Corporation, St. Louis, MO, U.S.A. unless otherwise indicated.

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2.2 Membrane Manufacturing Procedures The membrane manufacturing process followed here consists of three steps; (i) surface modification of ETFE mesh, (ii) preparation of proton exchange polymer and finally (iii) casting the proton exchange polymer onto the surface modified ETFE mesh to produce high conductive PEM. First, ETFE mesh was chlorinated by immersion in 15% sodium hypochlorite, into which phosphoric acid was carefully added until constant light bubbling was achieved. The solution was then stirred overnight. After chlorination, the mesh was rinsed with water and subsequently hydroxylated by placing in an aqueous 1M sodium hydroxide (NaOH) solution overnight. This method resulted in the preparation of a hydroxylated surface on the ETFE mesh rendering it adhesive or optional chemically reactive. The schematic of entire surface modification of ETFE mesh is shown in Figure 3. Second, the proton exchange polymer was prepared by adding the desired proportions of monomers acrylic acid, styrene, and vinyl sulfonic acid, the chemical structure of each of these monomers is shown in Figure 4, to a vial followed by a small amount of benzoyl peroxide to act as a radical initiator. For example (exact ratios are not mentioned here for proprietary reasons), 5 ml of acrylic acid, 8 ml of styrene, 0.1 ml of vinylsulfonic acid and 11.0 mg of benzoyl peroxide would form a typical membrane cocktail. Enough ethanol was then added to achieve a homogeneous solution. This solution was then very slowly heated in a sand bath to a temperature between 100oC to 110oC. After polymerization was complete, it was removed from the sand bath and allowed to cool at ambient temperature (25oC). After cooling, the polymer was isolated and then re-dissolved in ethanol to prepare the solution cocktail into which the hydroxylated ETFE mesh was placed. The treated (hydroxylated) ETFE mesh coated with the proton exchange polymer was then spun dry in a centrifuge. This casting procedure was repeated with drying in between each casting step. The new proton exchange polymer is based on styrene-acrylic acid-vinylsulfonate (SAS) and will subsequently be referred to as “SAS” polymer. The complete procedure to develop new SAS polymer exchange membrane (SASPEM) using the patented polymer surface modification technology [52-53] is shown briefly in Figure 5. The general structure of SASPEM indicated in Figure 5 wherein the SAS polymer matrix that exchanged protons, denoted by light green, was cast onto the mechanically stable hydroxylated ETFE mesh to form the final SASPEM. Since only the surface structure of the ETFE was modified through chemical processing, it was anticipated that the ETFE mesh would provide the structural support of the new SAS PEMs whereas the SAS proton exchange polymer would act as an efficient proton exchange media. Initial tests demonstrated (see Fig. 9) that increased sulfonate functionality (increased

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vinylsulfonic acid) to the polymer backbone provides increased proton conductivity by allowing the formation of proton migration channels throughout the polymer membrane. H H H

H H

H

H

H H At the Polymer Surface

NaOCl pH = 3 Cl Cl Cl

Cl Cl

Cl

Cl

Cl Cl At the Polymer Surface

NaOH

OHOH OH

OH OH

OH

OH

OH OH At the Polymer Surface

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Figure 3. Schematic of ETFE mesh functionalization processes chemically.

O O

O

O H

Vinyl Acetate

O

Styrene

S

H

Vinylsulfonic Acid

Figure 4. Monomer units used to form proton exchange SAS polymer. Polymer compositions formed from these species tend to be very hydrophilic and readily pass protons.

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Susanta K. Das H H H

H H

H

=Atoms = Functional Groups

H

H H At the Polymer Surface

BT Process- A

At the Polymer Surface

Paint or Glue

At the Polymer Surface

Process- B

Attach Other Species

At the Polymer Surface

At the Polymer Surface

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(a)

(b)

(c) Figure 5. (a) Schematic of the patented [52-53] polymer surface modification scheme, (b) general structure of proton exchange polymer cast onto ETFE mesh, and (c) final product of SAS proton exchange membrane.

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3. EVALUATION OF PROTON EXCHANGE MEMBRANE (PEM) PERFORMANCE Performance of proton exchange membrane is judged based on its ability for proton conductivity, resistance, water uptake, thermal and mechanical stability, swelling and fuel cross-over factors. Among these properties, conductivity and resistance are the primary metric for membrane performance.

3.1 Conductivity Conductivity is a bulk parameter which represents the ability of a material to conduct electricity. Thus it is independent of a membrane sample's dimensions. Low bulk membrane conductivity in a fuel cell results in a fuel cell with larger voltage losses which results in lower performance. Assuming all other components are equal, a fuel cell with high bulk membrane conductivity will have better performance due to smaller voltage losses. The units of conductivity are siemens per centimeter. A siemen is equal to 1/ohm. Conductivity is represented by the variable σ (sigma). The relation between conductivity and resistivity (denoted as ρ ) is σ = 1 ρ . Membrane conductivity can be the result of charge transfer from

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ions or electrons.

(a)

(b) Figure 6. Schematic of (a) in-plane and (b) through-plane conductivity test method.

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Susanta K. Das

The challenges in measuring membrane conductivity are in separating ionic creation from ionic conduction. When making a measurement of conductivity, if ionic creation is included, the measured value will be much lower than the actual bulk conductivity value. In general, it is assumed that the membrane conductivity is the result of charge transfer from ions. If a researcher suspects that the conductivity is a result of charge transfer from electrons, measurements of membrane conductivity must separate the voltage losses associated with ion creation from the voltage losses associated with ion conduction. When membranes are assembled into MEAs (Membrane Electrode Assemblies), several variables effect the conductivity measured, including inconsistency in the MEA assembly, contact resistance issues and flow field issues. Generally, in-plane and through-plane conductivity measurements are performed as shown in the Figure 6. In isotropic membranes, in-plane conductivity is the same as through-plane conductivity. In very anisotropic samples, the absolute values of conductivity will differ. But, the character of the in-plane curve (conductivity vs. RH) will be the same as the character of the through-plane curve. The advantage of measuring the character of the sample, even if the absolute conductivity value is not exact, is that the polymer chemist can quickly determine if the polymer is more or less dependent on water than Nafion®.

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3.1.1 Through-Plane Conductivity Measurements Through-plane conductivity is the measurement of interest in an operating fuel cell. However, to obtain a good through-plane conductivity measurement, the membrane must be assembled in a Membrane Electrode Assembly (MEA) with a high quality catalyst layer that is well-bonded to the membrane. A two-electrode measurement, using AC impedance or current interrupt, is used to obtain the conductivity measurement. In order to make accurate membrane conductivity measurements, the effects of the following variables must be considered: • Contact resistances between the diffusion media and the catalyst layer. • Contact resistances between the diffusion media and the current collector. • Diffusion media resistance. • Interactions within the catalyst layer. • Catalyst membrane interface. While through-plane conductivity is a more direct measurement of the membrane’s conductivity in a fuel cell, the equipment and resource costs are much higher than that required for in-plane conductivity measurements.

Advantages (i) (ii)

When the MEA includes a high quality, well-bonded catalyst, this is a very good measure of a membrane’s ability to conduct ions. This approach will also yield other fuel cell data.

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Disadvantages (i)

This approach takes time – building the MEA, building the fuel cell, and interpreting the data collected. (ii) This approach is expensive. The researcher will need a complete fuel cell testing system in order to obtain membrane conductivity data. (iii) The researcher will need to know how to build a high-quality fuel cell, and test it, in order to obtain accurate membrane conductivity data.

3.1.2 In-Plane Conductivity Measurements In-plane conductivity testing is a good solution because the issues related to inconsistency in the MEA assembly, contact resistance and flow field mentioned above can be easily eliminated. In addition, the sample is easier to get on test, as assembly in an MEA is not required, and equilibrates faster with no catalyst layer or GDL.

Advantages Assembly of the sample under test is much simpler. It is not assembled into an MEA. Instead, the pure membrane is clamped into the conductivity cell. (ii) This approach requires less equipments and supplies than that needed for fuel cell testing as required for through-plane, making it a less expensive approach. (iii) This method does not require an MEA. So, any variables introduced by the MEA are eliminated. (iv) When done properly, a four-electrode, in-plane measurement provides a more direct measurement of the conductivity of the membrane.

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(i)

Disadvantages (i)

This approach may be less effective on anisotropic membranes. However, the character of the membrane (conductivity vs. RH curve) will remain constant even in anisotropic samples. (ii) Some care must be taken to ensure that the measurement is measuring ion conduction and not electron conduction. (iii) In some cases, the measurement may be of surface conduction and not a true bulk characteristic. In this case, samples of varying thicknesses must be used to confirm that the measurement is a bulk measurement.

3.2 Resistance Resistivity ( ρ ) is a bulk parameter which represents a material's tendency to resist the flow of electrical charge. Resistance (R) is a material quantity. Thus it is dependent on a sample's dimensions. The unit of resistance is Ohms (Ω). Resistance is related to conductivity and resistivity through the following equation:

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Susanta K. Das

R=

(ρ ∗ L) = A

L (σ ∗ A)

(1)

where L is the length parallel to charge flow and A is the cross sectional area perpendicular to the charge flow. Thus membrane conductivity can be calculated by measuring membrane resistance for a given sample and then using the samples geometry to calculate conductivity as

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σ=

1 L = ρ (R ∗ A)

(2)

Various experimental techniques were used for measurement of membrane conductivity and resistance [54-55]. A Two-electrode or four-electrode probe for conductivity measurement and different techniques for resistance such as: (i) current interrupt (iR), (ii) AC resistance, (iii) electrochemical impedance spectroscopy (EIS), and (iv) high frequency resistance (HFR) are commonly used in practice. The experimental methods mentioned above are not accurate enough with each having advantages and disadvantages [1-2, 55] and also complicated to implement at the membrane manufacturing stage in the laboratory. Therefore, a simple equipment requirement and measurement technique supported by a rigorous theoretical model has to be developed to measure the membrane conductivity and resistance at the manufacturing level. The membrane conductivity and resistance are a particularly important measure of single fuel cell’s (or fuel cell stack) electrical performance since it quantifies internal cell losses. Thus, it is desirable to monitor membrane conductivity and resistance during membrane development and manufacturing of fuel cell stacks because ohmic losses generate waste heat, which must be removed from the fuel cell, resulting in a decrease in overall electrical efficiency. In addition, since the fuel cell current densities are quite high compared to other electrochemical processes, even a very low degradation of conductivity and low ohmic resistance (milliohm) have a significant effect on overall fuel cell efficiency [1-2]. The resistance offered to proton flow in a fuel cell is primarily due to the membrane [56]. Besides membrane, the electrode-membrane interface also has a contribution to the overall cell resistance. But compared to the membrane resistance to protons flow, the interface resistances are very negligible [1-2]. Hence, in this study we developed a correct theoretical framework eliminating ambiguities in previous model [57] to measure the conductivity and resistance of a proton exchange membrane based on proton flow, as it is the only ion that migrate from one electrode to the other, and also developed a cost effective simple equipment set-up for the accurate validation of the theoretical model by experiment.

3.3 Theoretical Model to Measure Membrane Conductivity and Resistance Based on Proton Flow The details of the theoretical model were published in [58] but for completeness of the chapter we have presented it here with permission from the publisher, Elsevier. We use the profiles of protons flow between two cells to measure the conductivity and resistance of

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proton exchange membrane. A schematic of protons flow profile through the proton exchange membrane in the conductivity cell (electrolyte phase) together with the concentration profiles of protons in the acid and water phases are presented in Figure 7. To minimize the complexity of the mathematical model, we assume that the protons concentration gradient can be approximated by a single-step linear difference between the concentration at the cells and interfaces. Therefore, the molar flux of protons in the acid cell towards the membrane can be expressed as the mass transfer coefficient in the acid, τ a , multiplied by the concentration gradient [59]:

N = τ a ∇Cac

(3)

⇒ N = τ am ( Ca − Cam )

(4)

where

∇Cac =

( Ca − Cam )

(5)

ξa

τ am = τ a ξ a denotes the modified mass transfer coefficient. Here we have considered solution resistance to protons flow in acid cell,

ξ a , for accurate calculation of resistance

which was neglected in [57]. We further assume that the resistance offered to protons flow by the proton exchange membrane can be modeled as resistor, with resistance R, which has a potential difference of Copyright © 2009. Nova Science Publishers, Incorporated. All rights reserved.

( Cam − Cmw )

across the membrane (see Fig. 7). We assume that the number of protons

leaving the acid cell per unit time equals the molar flux of protons through the membrane i.e. flow of protons satisfy the steady state condition. Thus, according to Ohm’s law, V = IR , we obtain:

NR = Cam − Cmw

(6)

where the molar flux (proton flow – similar to electron flow), N, is analogous to the current flow.

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Susanta K. Das

Figure 7. Schematic of concentration profiles of protons in the acid and water cells. Ca is the proton concentration in bulk of acid phase, Cw is the proton concentration in bulk of water phase, Cam is the proton concentration on the interface of acid and membrane phase, Cmw is the proton concentration on the interface of membrane and water phase, R is the resistance of membrane to proton flow. All the concentrations,

C 's , vary with time.

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Similarly for the water cell, we assume that there is no accumulation of protons in the membrane, i.e. the steady state condition is maintained, then we obtain:

N = τ w∇Cwc

(7)

⇒ N = τ mw ( Cmw − Cw )

(8)

Where

∇Cwc =

( Cmw − Cw ) ξw

(9)

τ w is the mass transfer coefficient in water phase, τ mw = τ w ξ w is the modified mass transfer coefficient between membrane-water interface and water phase, and

ξ w represents

solution resistance to protons flow in water cell which was neglected in [57]. Adding equations (4) and (8), we obtain

⎛ 1 1 ⎞ + Cam − Cmw = ( Ca − Cw ) − N ⎜ ⎟ ⎝ τ am τ mw ⎠

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where Ca and Cw are function of time. Substituting equation (10) into equation (6), we get,

Nr = ( Ca − Cw )

(11)

where

⎛ 1 1 ⎞ + r =⎜R+ ⎟ τ am τ mw ⎠ ⎝

(12)

denotes the total effective resistance offered by the membrane, the acid-membrane and the water-membrane interfaces. Since Ca = Ca (t ) , we can write the molar flux of proton concentration in the acid cell as

N =−

dCa dt

(13)

As the steady state condition is maintained throughout the systems, the rate of loss of protons from the acid cell must be equal to the rate of gain of protons in the water cell. Thus we obtain,

−

dCa dCw = dt dt

(14)

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Substituting equations (13) in equation (11) we get,

−r

dCa = Ca − C w dt

(15)

Using equation (14), equation (15) can be written as

⎛ dC ⎞ r ⎜ w ⎟ = Ca − C w ⎝ dt ⎠

(16)

Differentiating equation (16) once with respect to time, t, we get:

r

d 2Cw dCa dCw = − dt 2 dt dt

Using equation (14), equation (17) can be reduced as:

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(17)

16

Susanta K. Das

d 2Cw 2 dCw + =0 dt 2 r dt

(18)

The equation (18) represents a second order ordinary differential equation which can be solved easily [58] for the proton concentration in the water cell, Cw (t ) . Let

Cw (t ) = eqt

(19)

be a trail solution [13] of equation (18) rather than the one assumed in [57]. Using equation (19), the general solution of equation (18) can be obtained as [60]:

Cw (t ) = c1 + c2 e

⎛2⎞ − ⎜ ⎟t ⎝r⎠

(20)

where c1 and c2 are the arbitrary constant, which can be determined from the limiting case. Let the value of Cw ( t ) at time t = 0 be C0 and the value of Cw ( t ) at time t = ∞ be

C f . The values of C0 and C f can be determined from the experiment. Substituting values of Cw ( 0 ) and Cw ( ∞ ) in equation (20) we can find the values of c1 and c2 . Thus the final

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solution becomes

⎛ ⎡ C − C0 ⎤ −t ( r 2) ⎞ Cw (t ) = C f ⎜1 − ⎢ f e ⎟ ⎜ ⎢ C f ⎥⎥ ⎟ ⎣ ⎦ ⎝ ⎠

(21)

Equation (21) implies that the protons concentration in the water cell increases with time and attains its final value. The equation (21) can be written in a closed form as:

(

y = B 1 − Ae −t τ

)

(22)

where A , B are constants and τ = r 2 serves as a time constant. At time t = r 2

(23)

from equation (21), we obtain:

⎛ ⎡ C − C0 ⎤ ⎞ Cw (t ) = C f ⎜1 − 0.3679 ⎢ f ⎥ ⎟⎟ ⎜ C ⎢ ⎥⎦ ⎠ f ⎣ ⎝

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The equation (24) is correctly derived and eliminated numerical error in equation (23) of reference [57]. The values of C0 and C f can be determined using a simple experiment described in the following section. In equation (24), as the right hand side is completely known, the value of Cw (t ) can be easily found. Using equation (24), the profiles of protons flow in the water cell, Cw ( t ) , can be plotted as a function of time, t, from which the value of t corresponding to the calculated value of

Cw ( t ) can be determined and consequently the proton transfer rate through the membrane

can be easily obtained. Let the value of total time, t, be T. Hence, we obtain from equation (12) and (23):

⎛ 1 1 ⎞ + r =⎜R+ ⎟ = 2t = T τ am τ mw ⎠ ⎝

(25)

where r has the units of time. Physical interpretation of equation (24) along with equation

(25) is that to attain a particular value of Cw ( t ) i.e. a specific amount of protons

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concentration in the water cell, the total time required will be r. The greater the value of r, the longer time the membrane would take to allow those specific numbers of protons to pass into the water cell. It implies that r provides the measurement of the resistance of the membrane based on proton flow and r is directly proportional to the total time, T. Since the same experimental set-up and conditions are maintained in all the experiments to determine the values of C0 and C f , as required in equation (24), the value of ⎡⎣(1 τ am ) + (1 τ mw ) ⎤⎦ remains constant throughout. Hence the value of r depends only on R (according to equation (25)) i.e. the resistance of the membrane. In this simple way, we can measure the membrane conductivity and resistance offered to protons flow by each individual proton exchange membrane.

3.4 Experimental Set-Up and Procedures According to equation (24), to find the values of C0 and C f , which ultimately leads to the calculation of membrane conductivity and resistance, R, the following experimental set-up and procedures are followed.

3.4.1 Measurement of Proton Transfer Rate and Membrane Resistance As shown in Figure 8, the experimental set-up consists of two cells. The proton exchange membrane is placed in the PEM holder between the two cells. Silicone grease was applied around the outer edge of the PEM holder to ensure that transfer of protons from one cell to the other would not occur without passing through the PEM. First, same amount of de-ionized

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water was added to each cell. Then cells were allowed to equilibrate so that the water on each side of the membrane achieved the same depth. A pH meter was fitted into each cell and a monitor connection with pH meter was then established within each cell i.e. on both side of the membrane to accurately record the pH readings. The pH meter must be quite accurate to detect small changes in the pH of the two cells. As soon as the pH meters have reached in equilibrium, few drops of de-ionized water were added to the left cell and few drops (number of drops are same for both cases) of concentrated HCl solution were added to the right cell simultaneously. This procedure ensures the volume of liquid in each cell remained the same so that there would be no liquid forced through the membrane by pressure difference (i.e. avoided pressure driven flow). In this way we obtained, one acid cell containing concentrated HCl solution (20% HCl) and a water cell containing de-ionized water. The pH and temperature of both cells were recorded immediately. A half-reaction circuit with appropriate voltage was established between the two cells to maintain the electron-neutrality and brought the high proton concentration in the acid cell. Since there would be a negative gradient in the concentration of protons between the two cells, protons will move from the acid cell to the water cell. Thus, theoretically, protons must try to make their way across the membrane to migrate to the water cell. Therefore, pH readings of both acid cell and water cell were then acquired at a regular time interval for at least 90 minutes or longer if possible. All measurements were carried-out without stirring under quiescent conditions.

Electric circuit Acid cell

Water cell

PEM holder

Digital pH meter Figure 8. Membrane resistance test apparatus: two-cell method. Acid cell, water cell and test membrane sections are labeled and indicated by arrow sign. pH meter is attached into each of the cells and pH is

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recorded at a regular time interval through the digital pH meter. A half-reaction circuit is established between the cells.

3.4.2 pH Measuring Process and Calculation of Proton Concentration • • • • •

At the beginning of the experiment, the pH of both the acid cell and water cell were measured. Once the experiment was started, pH was measured at a regular interval of time. The experiment was terminated after the pH of both cells had reached equilibrium. The final value of pH in each cell (acid and water) was recorded. + − pH The concentration of protons was then obtained using the relation: ⎡⎣ H ⎤⎦ = 10 .

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3.4.3 Measurement of Proton Transfer To ensure protons are passing from acid cell to the water cell, initial proton conductivity tests were performed by measuring the pH at each cell for a long time period without placing any membrane. In principle, the pH of the acid cell should increase due to the loss of protons while the pH of the water cell should decrease due to receiving of protons. Initial test confirms the change in pH at each of the cells indicating that protons are migrated from acid cell to the water cell. Next, keeping the same initial concentration of protons in both the acid and water cells, commercial proton exchange membrane (PEM) Nafion® [4] and several prototypes in-house SAS (styrene-acrylic acid-vinylsulfonate) membranes were experimented using the same experimental set-up and experimental conditions. The membrane is placed between two-sided PEM holders as shown in figure 8. The pH and temperatures of both cells are recorded during entire experimentation at a regular time interval for a period of at least 90 minutes or longer wherever possible.

4. RESULTS AND DISCUSSIONS To validate the theoretical model to measure membrane resistance to protons flow, according to equation (24) to determine C0 and C f , we conducted experiment based on two-cell model developed in [58] and described in details in section 3. Initial proton transfer measurement is conducted using the apparatus set-up and procedures mentioned in section 3. Figure 9 represents the change of pH in water cell as a function of time without placing membrane between the two cells in order to test diffusion rate of protons (conduction of protons) at room temperature (25oC). Figure 9 shows a complete profile of change of pH in the water cell which clearly displays three distinct phases – an induction phase representing the time taken for the acid to diffuse into the de-ionized water to release protons up to the moment of time to start transfer of protons from acid cell to water cell i.e. total time taken to gain first proton into the water cell; a proton transfer phase wherein protons started passing rapidly into the water cell - lowering the pH of the water cell, and an equilibrium phase

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wherein the pH of the initially water-only cell was lowered approximately to that of the initially acidified cell. The proton transfer phase is used to determine the concentration of protons (i.e. total amount of protons transferred) into the water cell per minute and the time from initiation of proton transfer into the initially water-only cell till attainment of the equilibrium pH between the two cells. In Figure 9, the curve consists of three distinct segments for the rate of change of pH as a function of time: an initial slightly negative slope in induction phase, a greater negative slope in proton transfer phase, and finally a slightly negative slope in equilibrium phase. We divided the profile of pH obtained in Figure 9 into three separate curves, one for each region of different slope, and a linear regression line is fitted to each curve segment with corresponding linear equations. Let the regression lines be y1 = m1t + d1 (26)

y2 = m2t + d 2

(27)

y3 = m3t + d3

(28)

And respectively, in three phase regions. Here slopes, m1 , m2 , and m3 represent the rate of change of pH in water cell at each of the three distinct phases; induction phase, transfer phase and equilibrium phase, respectively. We can now calculate the concentration of protons at each of the phases via the relation of concentration of protons:

⎡⎣ H + ⎤⎦ =10-pH

(29) 7.0 y =m t+d

pH (in water cell)

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6.0

1

1

1

induction phase

5.0

transfer phase

y =m t+d 4.0

2

2

2

equilibrium phase

3.0

y =m t+d 3

3

3

2.0 1.0 0.0 0

profiles of rate of chnage in pH 10

20

30

time, t (min.)

40

50

Figure 9. Experimental results for the change of pH in water cell as a function of time with no membrane to determine the rate of diffusion of protons. Linear regression equations at each of the three distinct phases: induction phase, transfer phase and equilibrium phase, are shown in the onset. Slope of the curves indicate the rate of change of pH in each phase. This figure is reproduced from [58] with the permission from Elsevier Publications.

The strongly negative slope curve, in transfer phase, represents the maximum rate of proton transfer per minute and it can be obtained by using equations (27) and (29). The two

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intersections of the three curves given in equations (26)-(28) will provide the indication of transfer phase slope start and end time. Thus we can calculate slope start time, induction time, as:

t1 =

d 2 − d1 m1 − m2

(30)

and the slope end time, equilibrium time, as:

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t2 =

d3 − d 2 m2 − m3

(31)

To test the conductivity and resistance of the proton exchange membrane, we placed the membrane between the two cells as shown in figure 8 and followed the procedures discussed in section 3. Keeping the same initial concentration of protons in the acid and water cells, widely commercially used Nafion® [6] membrane and several prototypes of in-house SAS membranes were examined using the same experimental set-up and conditions. Figure 10 represents the rate of change of pH in water cell for different membranes tested here as a function of time at 80oC. As mentioned above, since we maintained same initial concentrations of protons in both of the cells at the time of starting experiment for each of the membrane case, the initial pH was same for each case as can be seen from Figure 10. In Figure 10, we can see that the variations of three distinct phases for each of membranes; induction phase, transfer phase and equilibrium phase, as a function of time. It is due to the resistance of each membrane offered to the protons flow while it passes through the membrane. We calculated the rate of change (slope) of pH profiles for each of the tested membrane, using equations (26)-(28), whose pH profiles in water cell is displayed in Figure 10. Figure 11 shows the rate of change of pH profiles in water cell at 80oC. In Figure 11, we can see that a steady constant slope at the induction phase, a significant slope variations in the transfer phase as protons started migration from acid cell to water cell to equilibrate the negative gradient in protons concentration created in the water cell, and finally a steady constant slope in the equilibrium phase which signals the attainment of pseudo-equilibrium in slopes between the acid cell and the water cell. These characteristics of slopes of pH profiles in water cell ascertain that how the proton conduction is taking place between the two cells. Figure 12 represents the concentration profiles of proton flow in the water cell at 80oC as a function of time for various membranes and the corresponding results obtained using the theoretical model, given in equation (24). For theoretical calculation, the values of C0 and

C f obtained experimentally through the rate of change of pH as shown in Figure 11 were used. In Figure 12 and all subsequent plots containing experimental and theoretical data, dotted, dashed and solid lines show the experimental and symbols represent the theoretical results respectively. There is an excellent agreement between the experimental and theoretical profiles of proton flow (see Figure 12) in all of the three distinct phases. The proton flow profile peak represents the maximum rate of proton transfer. Among the profiles of proton flow for different membranes, the SAS type I PEM had the highest peak and Nafion® 212 has the lowest peak. This indicates that the SAS type I PEM is able to transfer protons

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significantly faster than Nafion® 212 membrane at 80oC. Comparing results presented in Figure 12, we can see that both of the SAS type PEMs (SAS type I and SAS type II) is able to transfer higher number of protons per unit time than the commercial membrane Nafion® 212. 7.0

y = m t+d 1

1

induction phase

6.0

SAS type I SAS type II

pH in water cell

Nafion 212 y = m t+d

5.0

2

2

transfer phase

4.0 equilibrium phase y = m t+d 3

3

3.0

2.0 0

10

20

30

40

time, t (min.)

50

60

Figure 10. Experimental results for the change of pH in water cell as a function of time with different membranes at 80oC. Linear regression equations at each of the three distinct phases: induction phase, transfer phase and equilibrium phase, are shown in the onset. Slope of the curves indicate the rate of change of pH in each phase.

0.0

Slope (rate of change) of pH profile in water cell

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0.2

-0.2 SAS type I -0.4

SAS type II Nafion 212

-0.6

-0.8

-1.0

-1.2

-1.4 0

10

20

30

40

50

60

70

time, t (min.)

Figure 11. Slopes (rate of change) of pH profiles in water cell as a function of time with different membranes at 80oC. SAS membranes show very sharp decrease of slopes in the transfer phase as compared to Nafion® 212 membrane.

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Protons concentration in water cell (mole)

20

SAS type I [experiment]

16

SAS type I [theory - Eq. (24)] SAS type II [experiment] SAS type II [theory - Eq. (24)] Nafion 212 [experiment]

12

Nafion 212 [theory - Eq. (24)]

8

4

0 0

10

20

30

40

50

60

70

time, t (min.)

Figure 12. Concentration profiles of protons flow in water cell at 80oC as a function of time for different membranes. Symbols represent experimental results. Solid-, dashed- and dotted-lines represent theoretical model predictions for different membranes.

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4.1 Thermal Stability of Membrane Conductivity To confirm the stability of the SAS membranes under hydrolytic conditions i.e. fuel cells operating conditions, the membranes were immersed into water at different temperature for 24 hours and then re-evaluated for their proton transfer capacity. The proton transfer capacity for each of the membranes tested was then determined in the range of 25oC to 90oC in order to judge the membrane conductivity at different temperatures. The experimental procedures discussed in section 3 were followed to accomplish this. The standard deviation from the mean values of proton transfer capacity was ±2% (minimum 5 trials each). Figure 13 shows the maximum proton transfer capacity for each of the membrane examined at different temperatures. Maximum proton transfer capacity was determined using the highest peak slope as of Figure 11 at different temperatures. From Figure 13, it can be seen that both the SAS type membranes were able to transfer higher number of protons in the temperature range 25oC ~90oC than the Nafion® 212 membrane. Maximum protons transfer rate increased with increasing temperature for all of the membranes. The data presented in Figure 13 also indicates that the proton transfer rate is almost constant between 25oC~60oC. The rate of proton transfer for the SAS membrane increased rapidly after 60oC compared to Nafion® 212 membrane. At 80oC, SAS type I membrane provided the highest proton transfer rate, approximately 18 moles per minute, compared to Nafion® 212 membrane which exhibited a maximum proton transfer rate of approximately 1.8 moles per minute. This indicates that the SAS type I membrane has proton transfer capability approximately 10 times faster than Nafion® 212 membrane at 80oC. Figure 13, further indicates a slow decrease in proton transfer after 80oC for SAS type I membrane as temperature increases. Since the protons

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Susanta K. Das

present in the water cell are in the form of H3O+ and not simply H+, it is not known instantly what the significance of this shifted trend after 80oC when considering a hydrogen fuel source. Further experimental investigation is required to understand this trend.

Maximum protons transfer capacity (moles/min.)

20

16

SAS type I [experiment]

12

SAS type I [theory - Eq. (24)] SAS type II [experiment] SAS type II [theory - Eq. (24)] Nafion 212 [experiment]

8

Nafion 212 [theory - Eq. (24)]

4

0 20

30

40

50

60

70

80

90

100

o

Temperature ( C)

Figure 13. Maximum protons transfer capacity among membranes at different temperatures. Dashed, dotted and solid lines represent experimental results and symbols represent theoretical results as of Equation (24).

Average protons transfer capacity (moles/min.)

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12

SAS type I [experiment] SAS type I [theory - Eq. (24)]

10

SAS type II [experiment] SAS type II [theory - Eq. (24)] 8

Nafion 212 [experiment] Nafion 212 [theory - Eq. (24)]

6

4

2

0 20

30

40

50

60

70

80

90

100

o

Temperature ( C)

Figure 14. Average protons transfer capacity among membranes at different temperatures. Dashed, dotted and solid lines represent experimental results and symbols represent theoretical results as of Equation (24).

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Figure 14 presents the average proton transfer capacity for each of the membranes examined at different temperatures. Average proton transfer capacity was determined using the average slope calculated from the proton transfer phase profiles in water cell as shown in Figure 10 at 80oC. From Figure 14 we see that the average proton transfer rate of SAS type I membrane is 10.5 moles per minute compared to 1.5 moles per minute for Nafion® 212. This implies that the SAS type I membrane has 7 times faster average protons transfer rate than Nafion® 212 membrane under the test conditions utilized in this study. Figure 15 represents the minimum time required for the protons to pass through the membrane at different temperatures. The minimum time required for each membrane is determined by the difference of the time proton transfer into the water cell commenced and the time it took to transfer protons at its highest capacity (determined by the peak in the concentration profiles at different temperatures as shown in Figure 12 at 80oC). Comparing the results given in Figure 15 reveals that SAS type I membrane took 71% less time at 25oC whereas 87% less time at 80oC than Nafion® 212 membrane to transfer a mole of protons. Both experimental and theoretical results are in good agreement and indicate that the SAS type I membrane took the lowest time to transfer protons at the highest rate.

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Minimun time required for protons to pass through the membrane (min./mole)

1.0

0.8

0.6

0.4 SAS type I [experiment] SAS type I [theory - Eq. (24)] SAS type II [experiment]

0.2

SAS type II [theory - Eq. (24)] Nafion 212 [experiment] Nafion 212 [theory - Eq. (24)]

0.0 20

30

40

50

60

70

80

90

100

o

Temperature ( C)

Figure 15. Minimum time required for protons to pass through the membrane at different temperatures. Dotted, dashed and solid lines represent experimental results and symbols represent theoretical results as of Equations (30)-(31).

The data presented in Figure 16 provides the average time required for one mole of protons to pass through the membrane per minute at different temperatures. As before, both experimental and theoretical results are presented in Figure 16. The average time was calculated by using the starting and ending times of proton transfer in the transfer phase profiles at different temperatures as shown in Figure 12. Both SAS type membranes took less average time to transfer a mole of protons compared to Nafion® 212 membrane. At 25oC SAS type I membrane took 70% less average time whereas at 80oC it took 85% less average time than the peer Nafion® 212 membrane. This implies that both SAS type membranes are able to transfer protons at higher rates at less average time per mole than the Nafion® 212. Figure 17 shows the induction time required for protons to pass through the membrane at different

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Susanta K. Das

temperatures. Induction time is the time each membrane takes the very first proton to pass through the membrane and reach at the pH electrode (start of slope in the proton transfer phase) in water cell. The induction time is calculated by the relation given in equation (30). From Figure 17 it can be seen that at low temperature (i.e. 25oC), Nafion® 212 membrane has a higher induction time compared to SAS type membranes. The induction time decreases with increasing temperatures. At 80oC, induction time of Nafion® 212 is 3 times higher than SAS type I membrane. This indicates that SAS type I membranes are able to start proton transfer 3 times earlier than Nafion® 212 membrane at 80oC.

Average time required for protons to pass through the membrane (min./mole)

1.0

0.8

0.6

0.4 SAS type I [experiment] SAS type I [theory - Eq. (24)] SAS type II [experiment] 0.2

SAS type II [theory - Eq. (24)] Nafion 212 [experiment] Nafion 212 [theory - Eq. (24)]

0.0 20

30

40

50

60

o

70

80

90

100

Temperature ( C)

100

SAS type I 80

SAS type II Nafion 212

Induction time (minute)

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Figure 16. Average time required for protons to pass through the membrane at different temperatures. Dashed, dotted and solid lines represent experimental results and symbols represent theoretical results as of Equations (30)-(31).

60

40

20

0

20

30

40

50

60

70

80

90

100

o

Temperature ( C)

Figure 17. Induction time, equation (30), required for protons to pass through the membrane at different temperatures. Dotted, dashed and solid lines represent results for SAS type I, SAS type II and Nafion® 212 membrane respectively.

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4.2 Relative Resistance Figure 18 represents the minimum relative resistance among membranes at different temperatures. The relative resistance is calculated according to equation (25). According to the theoretical model [58], the membrane’s relative resistance is directly proportional to the total time taken by the membrane to allow a specific amount of protons to pass through it. Using the concentration profiles of protons in the water cell, as shown in Figure 12 at 80oC, the relative membrane resistance at different temperatures was evaluated. The minimum membrane resistance was calculated by the time each membrane took to transfer one mole of protons at its peak transfer rate at different temperatures. Comparing results presented in Figure 18, it can be seen that at 25oC, the relative resistance for Nafion® 212 is 72% higher and at 80oC the relative resistance is 87% more than the SAS type I membrane. The membrane resistance has a great impact on the performance of low temperature proton exchange membrane fuel cells since high membrane resistance causes a drop of fuel cell’s overall Ohmic voltage [1-2]. From Figure 18 it can be seen that among the membranes reported in this study, the SAS type I membrane had the lowest resistance and the Nafion® 212 membrane has the highest resistance at different temperatures. The average relative resistance among membranes is presented in Figure 19. Average relative resistance was calculated using the average time each membrane took at different temperatures to transfer specific amount of protons (one mole) through the transfer phases at an average proton transfer rate as shown in Figure 13 at 80oC. In Figure 19, it can be seen that the average relative resistance of Nafion® 212 membrane is much higher than the SAS type membranes at different temperatures. At 25oC, the average relative resistance of Nafion® 212 membrane is 65% higher and at 80oC the average relative resistance is about 80% higher than the SAS type I membrane. Since low resistance is a requirement for the enhancement of low temperature PEM fuel cells performance, SAS type I membrane shows a promise to perform better than Nafion® 212 membrane.

Membrane's minimum relative resistance

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2.0

1.6

1.2

0.8 SAS type I [experiment] SAS type I [theory - Eq. (25)] SAS type II [experiment]

0.4

SAS type II [theory - Eq. (25)] Nafion 212 [experiment] Nafion 212 [theory - Eq. (25)]

0.0 20

30

40

50

60

o

70

80

90

100

Temperature ( C)

Figure 18. Membrane's minimum relative resistance is shown at different temperatures. Dashed, dotted and solid lines represent experimental and symbols represent theoretical results as of Equation (25).

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Membrane's average relative resistance

2.0

1.6

1.2

0.8 SAS type I [experiment] SAS type I [theory - Eq. (25)] SAS type I [experiment] 0.4

SAS type II [theory - Eq. (25)] Nafion 212 [experiment] Nafion 212 [theory - Eq. (25)]

0.0 20

30

40

50

60

o

70

80

90

100

Temperature ( C)

Figure 19. Membrane's average relative resistance is shown at different temperatures. Dashed, dotted and solid lines represent experimental results and symbols represent theoretical results as of Equation (25).

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4.3 Water Uptake Content The swelling characteristics of the membranes were determined by the water uptake measurements. First, the membrane samples were placed in deionized water for at least 24h at room temperature. Then the membrane samples were removed from the water, excess water on the membrane surface removed with absorbent paper (tissue paper) and the samples immediately weighted on a microbalance. The membrane samples were dried at temperatures ranging from 25oC to 90oC. The dry weight was then quickly obtained at each temperature. This procedure was repeated at least five times until satisfactory reproducibility was obtained. The water uptake content was then calculated using the following relationship:

water uptake content, w ( % ) =

wwet − wdry wdry

× 100

(32)

where wdry and wwet are the weights of the dry and wet membrane samples, respectively. Figure 20 shows the results of membrane water uptake content calculated according to Equation (32). Dotted, dashed and solid lines represent the SAS type I, SAS type II and Nafion® 212 membrane respectively. From Figure 20, we see that the water uptake content for Nafion® 212 membrane is very high and increased almost linearly with increasing temperature as compared to SAS type membranes. This implies that to conduct proton transfer efficiently Nafion® 212 requires higher amount of water presence (i.e. higher

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humidity level) as compared to SAS type membranes. It appeared that SAS type membranes are capable of transferring protons efficiently at low water content i.e. at low humidity level than Nafion® 212. The dependence of the liquid water uptake on preceding dehydration conditions could have significant implications for the use of membrane in PEMFCs [1-2]. For example, in one common mode of fabrication of membrane-electrode assemblies (MEAs), the membrane and electrodes are hot-pressed together at higher temperature (e. g. at 120oC). During this process, all water is lost from the membrane and the subsequent operating temperature to which the membrane will be exposed (e.g. 80oC) could result in incomplete rehydration. If less water is taken up by the membrane thus treated, a decrease in the maximum attainable conductivity will occur since the conductivity depends strongly on membrane water content and hence on membrane’s relative humidity. Thus, the increased water uptake by the Nafion® 212 membrane following high temperature dehydration (by hotpressing during MEAs manufacture) could be a disadvantage in fuel cell operations in low humidity range, though it's reasonable to believe that Nafion® 212 membrane may be able to regain up to the required hydration level (80% or higher) quickly once soaked into water/vapor. To understand membrane water uptake in further details a rigorous relative humidity (RH) cycle measurement will be studied in the near future to verify the performance of these membranes in terms of conductivity, resistance, temperature, relative humidity and other parametric conditions. 35

SAS type I 30

SAS type II

Water uptake content

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Nafion 212 25

20

15

10

5 20

30

40

50

60

70

o

80

90

100

Temperature ( C)

Figure 20. Membrane’s water uptake content at different temperatures. Dotted, dashed and solid lines represent SAS type I, SAS type II and Nafion® 212 membrane respectively.

5. CONCLUSIONS The purpose of the current research was to develop a complete manufacturing procedure to fabricate a novel PEM, for fuel cells applications, that had the potential to be more

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efficient, robust and less expensive than Nafion®. The objectives of improved proton conductivity through a proton exchanging polymer matrix cast onto a very mechanically and chemically stable robust ETFE mesh were achieved by separating different PEM requirements and distributing them among different PEM polymers. In the present study, many of the structural and mechanical requirements were met by the ETFE mesh (not shown here for space limitation) while the proton exchange polymer media optimized proton transfer. Although the materials used to prepare these hybrid membranes were purchased from specialty chemical and polymer companies, the cost of ETFE used in the current study was approximately $353 m-2 and those of the other chemicals were under $20 m-2. Manufacturing costs are also relatively low and the preparation of the membrane cocktail employing the patented technology used in this study doesn’t require any expensive special equipments. It is thus anticipated that larger scale manufacturing of SAS type PEMs will be significantly less expensive than the manufacture of Nafion®-type membranes. The results of the current study indicate that the hybrid SAS type proton exchange polymer matrix may be tailored to meet a range of low temperature fuel cells requirements. In this study, we have shown that the basic requirements such as proton exchange capacities, relative resistance, temperature variations effect, induction time and water uptake content are significantly improved in SAS type membranes especially SAS type I membrane as compared to Nafion® 212 membrane. A detailed experimental set-up to validate the performance with the theoretical model has also been described. The results also show that the theoretical model predictions are in an excellent agreement with the experimental observations. These initial laboratory-based results reported here for the new SAS type PEMs are promising and further characterization of new SAS type PEM using industry-standard commercial equipment is presently underway which will be reported in future work. The results suggest that there is now a new route of fabricating cost-effective proton exchange membranes for fuel cell applications wherein one may focus more on the proton exchange capacity of the membrane allowing the structural properties of the membrane to be considered separately.

REFERENCES [1] J. Larminie, A. Dicks, Fuel cell systems explained, John Wiley & Sons, New York, 2000. [2] A. J. Bard, L. R. Faulkner, Electrochemical Methods: Fundamentals and Applications, John Wiley &Sons, New York, 2001. [3] US Department of Energy: Hydrogen, Fuel cells and Infrastructure Technologies Program (2006) (http://www1.eere.energy.gov/hydrogenandfuelcells/fuelcells). [4] Nafion®, E.I. DuPont de Nemours and Company, U.S.A. [5] H. L.Yeager and A. Steck, J. Electrochem. Soc., 128(9), 1880 (1981). [6] C. Yang, S. Srinivasan, A. B. Bocarsly, S. Tulyani and J. B. Benziger, Journal of Membrane Science, 237, 145 (2004). [7] M. Doyle, L. Wang, Z. Yang and S. K. Choi, J. Electrochem. Soc., 150 (11), D185 (2003). [8] K. D. Kreuer, J. Membr. Sci., 185, 29 (2001). [9] N. H. Jalani and R. Datta, Journal of Membrane Science, 264, 167-175 (2005). [10] R. Lawson, C. Wang, J. Hong, J. Ma, B. Fang and D. Chu, J. Electrochem. Soc., 154(1), B84 (2007).

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[11] T. A. Zawodzinski, J. Davy, J. Valerio and S. Gottesfeld, Electrochim, Acta, 40, 297 (1995). [12] L. J. M. J. Blomen and M. N. Mugerwa, Fuel cell systems, p. 614, Plenum Press, New York (1993). [13] S. Srinivasan, B. B. Dave, K. A. Murugesamoorthi, A. Parthasarathy and A. J. Appleby, Fuel cell systems, p. 37, L.J.M.J. Blomen, and M. N. Mugerwa, Editors, Plenum Press, New York (1993). [14] M. Doyle, M. E. Lewittes, M. G. Roelofs and S. A. Perusich, J. Phys. Chem. B, 105, 9387 (2001). [15] S. Chen, L. Krishnan, S. Srinivasan, J. Benziger and A. B. Bocarsly, Journal of Membrane Science, 243, 327 (2004). [16] G. Akovali and A. Oö zkan, Polymer, 27, 1277 (1986). [17] N. Carretta, V. Tricoli and F. Picchioni, J. Membr. Sci., 166, 189 (2000). [18] A.G. G. Garcia, P. N. Pintauro, M. W. Verbrugge and E. W. Schneider, J. Appl. Electrochem., 22, 204 (1992). [19] L. Hong and N. Chen, J. Polym. Sci., Part B, 38, 1530 (2000). [20] D. J. Jones and J. Rozie´re, J. Membr. Sci., 185, 41 (2001). [21] B. Gupta and G. G. Scherer, J. Appl. Polym. Sci., 50, 2129 (1993). [22] B. Gupta, F. N. Buechi, G. G. Scherer and A. Chapiro, J. Polym. Adv. Technol., 5, 493 (1994). [23] H. Kuhn, L. Gubler, T. J. Schmidt, G. G. Scherer, H. P. Brack, K. Simbek, T. Rager and F. Geiger, Proceedings of the 2nd European PEFC Forum, June 30-July 4, Lucerne, Switzerland, p. 69 (2003). [24] S. Hietala, M. Koel, E. Skou, M. Elomaa and F. Sundholm, J., Mater. Chem., 8, 1127 (1998). [25] S. Holmberg, J. H. Nasman and F. Sundholm, Polym. Adv. Technol., 9, 121 (1998). [26] S. Hietala, S. L. Maunu, F. Sundholm and T. Lehtinen, J. Polym. Sci., Part B, 37, 2893 (1999). [27] T. Lehtinen, G. Sundholm, F. Sundholm, P. Bjornbom, M. Bursell, Electrochim. Acta, 43, 1881 (1998). [28] D. I. Ostrovskii, L. M. Torell, M. Paronen, S. Hietala and F. Sundholm, Solid State Ionics, 97, 315 (1997). [29] S. Hietala, S. L. Maunu and F. Sundholm, J. Polym. Sci., Part B, 38, 3277 (2000). [30] M. Torkkeli, R. Serimaa, V. Etelaniemi, M. Toivola, K. Jokela, M. Paronen and F. Sundholm, J. Polym. Sci., Part B, 38, 1734 (2000). [31] C. Bailly, D. J. Williams, F. E. Karasz and W. MacKnight, J. Polymers, 28, 1009 (1987). [32] J. A. Kerres, J. Membr. Sci., 185, 3 (2001). [33] R. Nolte, K. Ledjeff, M. Bauer, R. Mu¨ lhaupt, J. Membr. Sci., 83, 211 (1993). [34] J. Kerres, W. Zhang, W. Cui, J. Polym. Sci. Part A, 36, 1441 (1998). [35] L. Kerres, W. Cui, M. Junginger, J. Membr. Sci., 139, 227 (1998). [36] H. Nakajima, S. Nomura, T. Sugimoto, S. Nishikawa and I. Honma, J. Electrochem. Soc., 149, A953 (2002). [37] P. Staiti, M. Minutoli and S. Hocevar, J. Power Sources, 90, 231 (2000). [38] J. M. Amarilla, R. M. Rojas, J. M. Rojo, M. J. Cubillo, A. Linares and J. L. Acosta, Solid State Ionics, 127, 133 (2000).

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[39] C. Poinsignon, I. Amodio, D. Foscallo, J. Y. Sanchez, Mater. Res. Soc. Symp. Proc., 548, 307 (2000). [40] S. P. Nunes, B. Ruffmann, E. Rikowski, S. Vetter and K. Richau, J. Membr. Sci., 203, 215 (2002). [41] B. Bonnet, D. J. Jones, J. Roziere, L. Tchicaya, G. Alberti, M. Casciola, L. Massinelli, B. Baner, A. Peraio and E. Ramunni, J. New Mater. Electrochem. Syst., 3, 87 (2000). [42] G. Alberti, M. Casciola and R. Palombari, J. Membr. Sci., 172, 233 (2000). [43] M. Rikukawa and K. Sanui, Prog. Polym. Sci., 25, 1463 (2000). [44] J. S. Wainright, M. H. Litt and R. F. Savinell, In Handbook of Fuel Cells; W. Vielstichm, A. Lamm, H. A. Gasteiger, Editors.; John Wiley & Sons Ltd.: New York, 3, 436 (2003). [45] R. J. Gillespie and E. A. Robinson, In Nonaqueous Solvent Systems; T. C. Waddington, Editor, p. 117, Academic Press, New York, (1965). [46] B. Xin and O. Savadogo, Electrochem. Commun., 2, 697 (2000). [47] K. T. Adjemian, S. Srinivasan, J. Benziger and A. B. Bocarsly, Journal of Power Sources, 109, 356 (2002). [48] M. Litt, R. Ameri, Y. Wang, R. Savinell and J. S. Wainwright, Mater. Res. Soc. Symp. Proc., 548, 313 (1999). [49] C. Hasiotis, Q. Li, V. Deimede, J. K. Kallitsis, C. G. Kontoyannis and N. J. Bjerrum, J. Electrochem. Soc., 148, A513 (2001). [50] L. Jorissen, V. Gogel, J. Kerres and J. Garche, J. Power Sources, 105, 267 (2002). [51] P. Choi, N. H. Jalani, T. M. Thampan and R. Datta, Journal of Polymer Science Part B: Polymer Physics, 44, 2183 (2006). [52] L. G. Beholz, Process for Producing Painted Polymeric Articles, US Patent No. 6,100,343 (2000). [53] L. G. Beholz, Apparatus for Treating Polymeric Material to Improve Surface Adhesion, US Patent No. 7,022,291 (2006). [54] Achieving Accurate and Reliable Resistance Measurements in Low Power and Low Voltage Applications, Keithley Instruments white paper, 2004. [55] M. Smith, K. Cooper, D. Johnson, L. Scribner, Fuel Cell Magazine, April/May issue, Webcom Communications Corp., 2005. [56] T. A. Zawodzinski, T. E. Springer, F. A. Uribe, S. Gottesfeld, J. Solid State Ionics, 60 (1993) 199. [57] B.V. Babu, N. Nair, J. Energy Edu. Sci. Tech. 13 (2004) 13-20. [58] S. K. Das and K. J. Berry, J. Power Sources, 173, 909 (2007). [59] R. Treybal, Mass Transfer Operations, McGraw-Hill, Singapore, 1980. [60] H. C. Edwards, D. E. Penney, Differential Equations: Computing and Modeling, 3rd ed., Prentice Hall, NJ, 2003.

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Chapter 2

THE DEVELOPMENT OF BIPOLAR PLATE MATERIALS FOR POLYMER ELECTROLYTE MEMBRANE FUEL CELLS (PEMFC) Yan Wang and Derek O. Northwood* Department of Mechanical, Automotive, and Materials Engineering, University of Windsor, 401 Sunset Avenue, Windsor, Ontario, Canada N9B 3P4

ABSTRACT

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With escalating oil prices and heightened environmental concerns, increasing attention is being paid to fuel cell technology. Polymer Electrolyte Membrane Fuel Cells (PEMFCs) have significant advantages compared with other types of fuel cell. PEMFCs have a solid electrolyte which provides excellent resistance to gas crossover. PEMFCs are considered as promising candidates for transportation applications because of their lower temperature operation and fast start-up. However, there are a number of barriers, e.g. price, weight and volume, which have prevented PEMFCs from being more widely used. Bipolar plates (BPs) are very important components in the PEMFC stack. They are designed to accomplish many functions, including: distribute the fuel and oxidant in the stack; facilitate water management within the cell; separate the individual cells in the stack; carry current away from the cell; and facilitate heat management. Currently, BPs are made of non-porous graphite because of its chemical and thermal stability in the PEMFC working conditions. However, their high price and the difficulties in machining the gas flow channels present problems for the application of non-porous graphite BPs. Also, non-porous graphite is very brittle, and is therefore readily broken, with the result that H2 and O2 could leak into the whole fuel cell system, which is potentially dangerous for fuel cell operation. Increasing attention is, therefore, being paid to the use of metallic and composite materials as replacements for non-porous graphite in BPs for PEMFCs. The ideal BP material should demonstrate high values of electrical conductivity, thermal conductivity, corrosion resistance and compressive strength and low values of gas permeability and density. In this paper, the three types of BP materials, i.e. non-porous *

Corresponding author. Tel.: +1 519 253 3000x4785; Fax: +1 519 973 7007. E-mail address: [email protected] (D.O.Northwood) [email protected] (Y.Wang)

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Yan Wang and Derek O. Northwood graphite, metallic and composite materials, are reviewed for their physical, chemical and electrochemical properties. For the metallic BPs, a wide variety of coating techniques and materials have been used in order to increase the corrosion resistance of the base materials. For the composite materials, different composite material compositions and processing methods have been used to increase the electrical conductivity and production rates. The most promising candidate materials for BPs that meet the design requirements are identified.

Keywords: PEMFCs, Bipolar plates (BPs), metallic, composite

1. FUEL CELLS Fuel cells are electrochemical devices which are continuously fueled by fuel and oxygen. Each cell consists of an electrolyte (a conducting ionic membrane) with an integrated porous anode and cathode. Hydrogen and/or hydrocarbon fuels react at the anode side, while oxygen (from air) reacts at the cathode side. The output is electrical energy in the form of direct current [1]. Although there are different types of fuel cell, generally a fuel cell assembly is composed of the following components [1]: • •

•

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•

Electrolyte membrane (with high ionic conductivity and retained hermetic seal) Anode (fuel side of the electrode with a porous composite structure containing ionconducting media, high surface areas for reaction, and catalysts lodged at the triple points where the reaction occurs among the gases, ions, and electrons) Cathode (air side of the electrode, with the same properties as the anode but in an oxidizing environment), and Separator and gas-flow structures (supplying air and fuel flow to the electrodes and also separating each cell hermetically while conducting electrons).

Fuel cells are environmentally friendly devices for energy conversion and power generation, and are one of the most promising candidates as a zero-emission power source. Hence, they are often regarded as one of the advanced energy technologies of the future. In reality, fuel cells are one of the oldest energy conversion devices known to humankind, although their development and deployment for practical applications lags far behind other competitive technologies, mainly heat engines such as the steam turbine and the internal combustion engine [2]. Nowadays, with escalating oil prices and increasing environmental concerns, increasing attention is being paid to the development of fuel cells [3-11]. The solidstate fuel cell has the potential to generate another new age, in the areas of distributed energy, a cleaner environment, and more efficient use of the earth’s natural resources. The need for cleaner energy-producing equipment is becoming a commercial necessity and will continue to grow in importance as the demand for oil outpaces production capacity in the next decade [12]. Advancements in materials science, chemical engineering, as well as in mechanical engineering, will be needed to ensure the chemical and mechanical long-term reliability of fuel cell technologies

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2. POLYMER ELECTROLYTE MEMBRANE FUEL CELLS (PEMFCS) 2.1 What is PEMFCs?

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Among the various fuel cell systems known today, the polymer electrolyte membrane fuel cell (PEMFC) system has proven to be an attractive and more promising option than alkaline, solid oxide, or molten carbonate fuel cell systems for power generation in portable, stationary, and mobile (automobile) applications [6]. Polymer electrolyte membrane fuel cells are also known as ion exchange membrane fuel cells (IEMFCs), solid polymer (electrolyte) fuel cells (SP(E)FCs), and proton exchange membrane fuel cells (PEMFCs). The use of a solid polymer membrane as an electrolyte separator in fuel cell applications offers several advantages, such as selectivity, system simplicity, and improved reliability compared with systems based on a liquid electrolyte. A schematic diagram of a PEMFC is shown in Fig 2.1. A PEMFC’s operating temperature is between 70-100°C. The electrolyte is a polymer membrane. The Nafion membrane, made by DuPont, is being widely used as an electrolyte in PEMFCs because it has a high conductivity and low equivalent weight. Platinum and platinum alloys supported on carbon black are being used as both anode and cathode electrocatalysts. The anode, cathode and overall cell reactions are given in equations 2.1, 2.2 and 2.3, respectively. Anode Reactions: 2H2 => 4H+ + 4e-

2.1

Cathode Reactions: O2 + 4H+ + 4e- => 2H2O

2.2

Overall Cell Reactions: 2H2 + O2 => 2H2O

2.3

 

Figure 2.1. A schematic illustration of a polymer electrolyte membrane fuel cell [13]

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2.2 Why Are PEMFCs so Important? PEMFCs have significant advantages compared with other type’s fuel cells. Because PEMFCs have a solid electrolyte, it provides excellent resistance to gas crossover. PEMFCs are operated in low temperature; therefore they are allow rapid start-up. Test results have shown that PEMFCs are capable of high current densities of over 2kW/l and 2W/cm2 [14]. These characteristics allow PEMFCs to be used in transportation, mobile/small devices and telecommunications. Figures 2.2 and 2.3 [15] show two examples where PEMFCs are used.

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Figure 2.2. Ballard laptop power [15]

Figure 2.3. Portable fuel cell power unit [15]

Nowadays, almost all large automotive companies are pursuing research work on PEMFC cars. Figs 2.4 to 2.7 [16] show examples of vehicles powered by PEMFCs. PEMFCs are the most promising fuel cells for use in the automotive industry because of their characteristics. All the major automotive companies are pursuing research work on PEMFC vehicles in order to take a lead position in future automotive markets.

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Figure 2.4. Ford Focus FCV [16]

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Figure 2.5. GM: Sequel [16]

Figure 2.6. Honda: FCX-V4 [16]

Figure 2.7. Toyota: FCHV-4 [16]

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2.3 How do PEMFCs Work?

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PEMFCs consist of a polymeric membrane as an electrolyte. Hydrogen or hydrocarbon fuels react at the anode side, while oxygen reacts at the cathode side. The output is electrical energy in the form of direct current. When hydrogen is used as the fuel, the final exhaust product is simple water. Fig 2.8 [17] illustrates how a PEMFC works. First, hydrogen or a hydrocarbon is reduced at the anode, and then hydrogen ions migrate to the cathode through the membrane. Oxygen from the cathode reacts with hydrogen ions and electrons to produce water. The theoretical potential for a PEMFC is 1.229V. However, because of activation polarization, ohmic polarization and concentration polarization (Fig 2.9), the working potential for PEMFC is about 0.7V. From Fig 2.9, we can see that activation polarization is the dominant type of polarization when the current density is low. With increasing current density, ohmic polarization becomes the main polarization mechanism. When the current density is very high, the dominant polarization mechanism is concentration polarization.

Figure 2.8. PEMFC working principle [17]

Figure 2.9. Ideal and actual fuel cell voltage/current characteristics [18]

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2.3.1 Activation Polarization Some electrode reactions are inherently slow and it is the slowness that gives rise to activation polarization. As with all chemical reactions, slow kinetics are associated with an activation energy. The same is true for electrode reactions. The extra voltage, which is called the activation overvoltage (ηact), is required for the electrochemical reactions [18, 19, 20]. This activation overvoltage is given by equation 2.4.

η act =

RT i ln anF io

2.4

where α is the electron transfer coefficient of the reaction at the electrode being addressed, i0 is the exchange current density, n is the number of electrons transferred in the reaction, and F is Faraday constant.

2.3.2 Ohmic Polarization Ohmic polarization is attributed to the resistance both to the electron transfer through the cell materials including the electrode, and to the ion transfer through the electrolyte. The ohmic overpotential can be calculated as follows [18, 20]:

η ohm = iR

2.5

where i is the current density and R is the total cell resistance (electronic, ionic, and contact resistance).

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2.3.3 Concentration Polarization A concentration gradient formed due to the depletion of the fuel at the electrode causes this type of polarization. The concentration polarization can be determined using the following equation [18, 20]:

η conc =

RT i ln(1 − ) anF iL

2.6

where α is the electron transfer coefficient of the reaction at the electrode, iL is the limiting current density, n is the number of electrons transferred in the reaction, and F is the Faraday constant.

2.3.4 Reactant Crossover and Internal Current Losses Fuel crossover is one of the major drawbacks of the most PEMFC systems. Fuel migration through the electrolyte to the cathode can cause a decrease in the cathode performance, and thus, a loss in overall fuel cell efficiency [20, 21].

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3. BIPOLAR PLATES AND THEIR DESIGN 3.1 What Are the Bipolar Plates? Bipolar plates are one of the most important components in PEMFCs. They are designed to accomplish many functions, such as, distribute the fuel and oxidant in the stack, facilitate water management with the cell, separate the individual cells in the stack and carry current away from the cell, and facilitate heat management [22]. Furthermore, the plates must be made of inexpensive, lightweight materials and must be easily and inexpensively manufactured. Efforts are underway to develop bipolar plates that satisfy these demands.

3.2 Why Are the Bipolar Plates Important?

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The present cost of fuel cell is about $200kW-1 which is the major barrier for commercialization in automotive applications [23]. Fig 3.1 and 3.2 [24] show us the weight and cost structure of PEMFCs. It can be seen that bipolar plates account for about 80% of the total weight and 45% of stack cost. In order to compete with an internal combustion engine, the cost of the total fuel cell stack should be reduced to $50 kW-1 and the cost of bipolar plates should be reduced to $10 kW-1. However, currently in the PEMFC stack, the cost of bipolar plates is about $90 kW-1. This is the incentive for research on less expensive materials and processes for the manufacture of bipolar plates for PEMFCs.

Figure 3.1. Weight structure of PEMFC stack [24]

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Figure 3.2. Cost structure of PEMFC stack [24]

3.3 Bipolar Plate Design

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Bipolar plates are designed in many configurations of flow fields in order to fulfill their functions. Currently, the “popular” designs of bipolar plates include the pin-type flow field, the series-parallel flow field, the serpentine flow field, the integrated flow field, the interdigitated flow field, and the flow-field designs made from metal sheets.

Pin-Type Flow Field Fig 3.3 [25, 26] shows the pin-type flow field for bipolar plates design. In the pin-type flow field, different shaped pins are used to form channels and the fuel and oxidant flow across the plates through the grooves formed by the pins. The advantage of this design is the low reactant pressure drop. However, the problem of this kind of design is that gases always flow through the least resistance areas, which leads to the formation of stagnant areas, uneven distribution of gases, inadequate product water removal and poor fuel cell performance [27].

Outlet

Inlet

Figure 3.3. Pin-type flow field [25]

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Straight Flow Field Fig 3.4[28] illustrates the straight flow design of bipolar plate. In this design, the gas flow-field plate includes a number of separated parallel flow channels connected to the gas inlet and exhaust headers, which are parallel to the edges of the plate. The problem of this type of design has been a low and unstable cell voltage after a period of operation because of cathode gas flow distribution and cell water management. Another problem is that the straight and parallel channels in the bipolar plates tend to be relatively short and have no directional changes. Therefore, the reactant gas has a very small pressure drop, which results in nonuniform flow distribution of reactant gas in the fuel cell stack [27].

Figure 3.4. Straight flow field [28]

Serpentine Flow Field Fig 3.5 [29] illustrates one type of serpentine flow field design for bipolar plates. A serpentine flow field can solve flooding due to inadequate water removal from the cell. Such a single serpentine flow field forces the reactant flow to traverse the entire active area of the corresponding electrode thereby eliminating areas of stagnant flow. However, this channel layout results in a relatively long reactant flow path, hence a substantial pressure drop and significant concentration gradients from the flow inlet to outlet [27].

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Figure 3.5. Serpentine flow field [29]

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Integrated Flow Field Fig 3.6 [30] shows one kind of integrated flow fields. In the integrated design, reactant gas flow field and cooling flow field are on the same plate surface. In this design, a fluid flow field plate is divided into a multiple of fluid flow sub-plates. Each sub-plate is electrically insulated from all other sub-plates of the same plate assembly, and has its own reactant flow field adjacent to the electrochemically active area of the nearby MEA. A cooling flow field may be positioned in-between and around each of the gas flow sub-plates. However, this kind of design can not maintain a uniform temperature distribution in the fuel cell stack [27, 30].

Figure 3.6. Intergrated flow field [30]

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Interdigitated Flow Field Fig 3.7 [31] shows one type of interdigitated flow fields. In an interdigitated flow field design, the dominant reactant flow is in the direction parallel to the electrode surface, and the reactant flow to the catalyst layer, required for electrochemical reaction and electric power generation, is predominantly by molecular diffusion through the electrode backing layer. Not only is molecular diffusion a slow process, which easily leads to the occurrence of large concentration gradients across the backing layer and a mass transfer limitation phenomenon for the cell operation, but it is also difficult to remove the water which exists in the porous region of the backing layer. This difficulty is compounded by the fact that the typical flow in the flow channels is laminar due to the small gas velocity and the small flow channel dimensions [27].

Figure 3.7. Interdigited flow field [31]

Flow-Field Designs Made From Metal Sheets Fig 3.8 [32] shows one kind of serpentine flow fields. Flow through porous carbon has also been proposed for improved water management; a better method may be the use of flow through porous metallic meshes (with high resistance to

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corrosion) to improve gas distribution on the cell plane. The plates comprise corrosionresistant thin metal sheets and reactant gas flow fields on the two outside surfaces of the sheets. Such a bipolar plate design eliminates the need for a separate cooling plate, decreases material usage for stack construction and reduces the weight and volume of the stack [27].

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Figure 3.8. Configurations for flow fields made by metal sheets [32]

3.4 Materials for Bipolar Plates Bipolar plates are required to perform many functions as has been discussed in section 3.1. In order to satisfactorily perform these functions, certain physical and chemical properties are required of the bipolar plate material. Table 3.1 provides a summary of the functions together with the required physical and chemical properties specific to that function. Table 3.1. Functions and required physical and chemical properties of bipolar plates [22] Function Distribution and management of fuel and oxidants and residual gases and liquids Conduct electrical current Facilitate heat management Separate the individual cells in the stack

Physical and chemical properties H2 permeability (dry, non-porous plates), bubble pressure (wet, porous plates), corrosion resistance Electrical conductivity Thermal conductivity Compressive strength

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Other physical properties of bipolar plate materials that are of importance to the performance of the fuel cell are coefficient of thermal expansion, density and hydrophobicity. Therefore, for a material to qualify for use as a bipolar plate, the criteria listed in Table 3.2 have been suggested [11, 33, 34]: Table 3.2. The criteria for the bipolar plates [11, 33, 34] Electrical conductivity Thermal conductivity Hydrogen/gas permeability Corrosion resistance Compressive strength Density Cost

Plate resistance 0) and condensation (m& phase < 0) is assumed, so that the mass balance equations for both phases are;

∇ ⋅ ((1 − sat )ρ g εu g ) = m& phase

(9)

and

∇ ⋅ (sat.ρ l εu l ) = m& phase

(10)

The momentum equation for the gas phase reduces to Darcy’s law, which is, however, based on the relative permeability for the gas phase (KP ) . The relative permeability accounts for the reduction in pore space available for one phase due to the existence of the second phase. The relative permeability for the gas phase is given by;

KPg = (1 − sat )KP

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137

and for the liquid phase;

KPl = sat.KP

(12)

The momentum equation for the gas phase inside the gas diffusion layer becomes;

u g = −(1 − sat )

Kp

μg

∇P

(13)

Two liquid water transport mechanisms are considered; shear, which drags the liquid phase along with the gas phase in the direction of the pressure gradient, and capillary forces, which drive liquid water from high to low saturation regions [19]. Starting from Darcy’s law, the following equation can write;

ul = −

Kp l

μl

∇Pl

(14)

where the liquid water pressure stems from the gas-phase pressure and the capillary pressure according to [19];

∇Pl = ∇P − ∇Pc = ∇P −

∂Pc ∇sat ∂sat

(15)

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Introducing this expression into Equation (14) yields a liquid water velocity field equation;

ul = −

KPl

μl

∇P +

KPl ∂Pc ∇sat μ l ∂sat

(16)

The functional variation of capillary pressure with saturation is prescribed following Leverett [19] who has shown that;

⎛ ε ⎞ Pc = σ ⎜ ⎟ ⎝ KP ⎠

12

f (sat )

(17)

f (sat ) = 1.417(1 − sat ) − 2.12(1 − sat ) + 1.263(1 − sat ) (18) 2

3

The correlation for interfacial liquid/gas tension can be expressed as [20]; Z

⎛ T −T ⎞ ⎟⎟ (1 × 10 − 3 ) σ = σ 1 ⎜⎜ c − T T 1 ⎠ ⎝ c

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where

σ 1 = 71.97

Tc = 647.35

T1 = 298.15

Z = 0.8105 .

The liquid phase consists of pure water, while the gas phase has multi components. The transport of each species in the gas phase is governed by a general convection-diffusion equation in conjunction which the Stefan-Maxwell equations to account for multi species diffusion, with the addition of a source term accounting for phase change; N ⎡ ∇P ⎤ ⎤ ∇M ⎞ M ⎡⎛ +⎥ ⎜ ∇y j + y j ⎟ + (x j − y j ) ⎢ − (1 − sat )ρ g εy i ∑ Dij ⎢ M j ⎣⎝ M ⎠ P ⎥⎦ ⎥ j =1 ⎢ ∇⋅ = m& phase ⎢ ⎥ ∇ T (1 − sat )ρ g εy i ⋅ u g + εDiT ⎢ ⎥ T ⎦ ⎣

(20)

where the subscript i denotes oxygen at the cathode side and hydrogen at the anode side, and j is water vapour in both cases. Nitrogen is the third species at the cathode side. In order to account for geometric constraints of the porous media, the diffusivities are corrected using the Bruggemann correction formula [21].

Dijeff = Dij × ε 1.5

(21)

The heat transfer in the gas diffusion layers is governed by the energy equation as follows;

∇ ⋅ ((1 − sat )(ρ g εCp g u g T − k eff , g ε∇T )) = εβ (Tsolid − T ) − εm& phase ΔH evap

(22)

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where the term ( εβ (Tsolid − T ) ), on the right hand side, accounts for the heat exchange to and from the solid matrix of the GDL. The gas phase and the liquid phase are assumed to be in thermodynamic equilibrium, i.e., the liquid water and the gas phase are at the same temperature. The enthalpy of evaporation for water is calculated as follow [20]; Z

ΔH evap

⎛ T −T ⎞ ⎟⎟ × 4186.80 = ΔH evap ,l ⎜⎜ c − T T 1 ⎠ ⎝ c

where ΔH evap ,l = 538.7

Tc = 647.35

T1 = 373.15

(23)

Z = 0.38 .

The potential distribution in the gas diffusion layers is governed by;

∇ ⋅ (λ e ∇φGDL ) = 0

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Implementation of Phase Change In order to account for the magnitude of phase change inside the GDL, expressions are required to relate the level of over- and undersaturation as well as the amount of liquid water present to the amount of water undergoing phase change. In the case of evaporation, such relations must be dependent on (i) the level of undersaturation of the gas phase in each control volume and on (ii) the surface area of the liquid water in the control volume. The surface area can be assumed proportional to the volume fraction of the liquid water in each cell. A plausible choice for the shape of the liquid water is droplets, especially since the catalyst area is Teflonated [19]. The evaporation rate of a droplet in a convective stream depends on the rate of undersaturation, the surface area of the liquid droplet, and a (diffusivity dependent) masstransfer coefficient. The mass flux of water undergoing evaporation in each control volume can be represented by [19];

m& evap = M H 2O ϖ N D k xm π Ddrop The bulk concentration

x w∞

x w 0 − x w∞ 1 − x w0

(25)

is known by solving the continuity equation of water vapor.

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to obtain the concentration at the surface x w0 , it is reasonable to assume thermodynamic equilibrium between the liquid phase and the gas phase at the interface, i.e., the relative humidity of the gas in the immediate vicinity of the liquid is 100%. Under that condition, the surface concentration can be calculated based on the saturation pressure and is only a function of temperature. The heat-transfer coefficient for convection around a sphere is well established, and by invoking the analogy between convective heat and mass transfer, the following mass-transfer coefficient is obtained [19];

k xm

c air D H 2O ⎡ ⎛D v ρ ⎢2 + 0.6⎜ drop ∞ g = ⎜ Ddrop ⎢ μg ⎝ ⎣

⎞ ⎟ ⎟ ⎠

12

⎛ μg ⎜ ⎜ ρ g DH O 2 ⎝

⎞ ⎟ ⎟ ⎠

13

⎤ ⎥ ⎥ ⎦

(26)

It is further assumed that all droplets have a specified diameter D drop , and the number of droplets in each control volume is found by dividing the total volume of the liquid phase in each control volume by the volume of one droplet;

ND =

sat.Vcv 1 3 πDdrop 6

(27)

In the case when the calculated relative humidity in a control volume exceeds 100%, condensation occurs and the evaporation term is switched off.

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The case of condensation is more complex, because it can occur on every solid surface area, but the rate of condensation can be different when it takes place on a wetted surface. In addition, the overall surface area in each control volume available for condensation shrinks with an increasing amount of liquid water present. Berning and Djilali [19] assumed that the rate of condensation depends only on the level of oversaturation of the gas phase multiplied by a condensation constant. Thus, the mass flux of water undergoing condensation in each control volume can be represented by;

m& cond = ϖ C

x w 0 − x w∞ 1 − x w0

(28)

2.2.3. Catalyst Layers The catalyst layer is treated as a thin interface, where sink and source terms for the reactants are implemented. Due to the infinitesimal thickness, the source terms are actually implemented in the last grid cell of the porous medium. At the cathode side, the sink term for oxygen is given by;

M O2

S O2 = −

4F

ic

(29)

Whereas the sink term for hydrogen is specified as;

S H2 = −

M H2 2F

ia

(30)

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The production of water is modelled as a source terms, and hence can be written as;

S H 2O =

M H 2O 2F

ic

(31)

The generation of heat in the cell is due to entropy changes as well as irreversibilities due to the activation overpotential [22];

⎤ ⎡ T (− Δs ) q& = ⎢ + ηact ,c ⎥ ic ⎦ ⎣ ne F

(32)

The local current density distribution in the catalyst layers is modelled by the ButlerVolmer equation [18-22];

⎛ CO i c = ioref,c ⎜ ref2 ⎜ CO ⎝ 2

⎞⎡ ⎛ α a F ⎞ ⎛ α F ⎞⎤ ⎟ exp⎜ η act ,c ⎟ + exp⎜ − c η act ,c ⎟⎥ ⎢ ⎟ ⎣ ⎝ RT ⎠ ⎝ RT ⎠⎦ ⎠

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ia = i

ref o ,a

⎛ CH2 ⎜ ⎜ C ref ⎝ H2

141

12

⎞ ⎡ ⎛ αa F ⎞ ⎛ α F ⎞⎤ ⎟ exp⎜ η act ,a ⎟ + exp⎜ − c η act ,a ⎟⎥ ⎢ ⎟ ⎣ ⎝ RT ⎠ ⎝ RT ⎠⎦ ⎠

(34)

2.2.4. Membrane The balance between the electro-osmotic drag of water from anode to cathode and back diffusion from cathode to anode yields the net water flux through the membrane [23];

N W = n d M H 2O

i − ∇ ⋅ (ρDW ∇cW ) F

(35)

The water diffusivity in the polymer can be calculated as follow [24];

1 ⎞⎤ ⎡ ⎛ 1 − ⎟⎥ DW = 1.3 × 10 −10 exp ⎢2416⎜ ⎝ 303 T ⎠⎦ ⎣

(36)

The variable cW represents the number of water molecules per sulfonic acid group (i.e. mol H 2O equivalent SO3−1 ).The water content in the electrolyte phase is related to water vapor activity via [25, 26];

cW = 0.043 + 17.81a − 39.85a 2 + 36.0a 3 cW = 14.0 + 1.4(a − 1)

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cW = 16.8

(0 < a ≤ 1) (1 < a ≤ 3) (a ≥ 3)

(37)

The water vapor activity given by;

a=

xW P Psat

(38)

Heat transfer in the membrane is governed by;

∇ ⋅ (kmem ⋅ ∇T ) = 0

(39)

The potential loss in the membrane is due to resistance to proton transport across membrane, and is governed by;

∇ ⋅ (λ m ∇φ mem ) = 0

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2.2.5. Cell Potential Useful work (electrical energy) is obtained from a fuel cell only when a current is drawn, but the actual cell potential, Ecell, is decreased from its equilibrium thermodynamic potential, E, because of irreversible losses. The various irreversible loss mechanisms which are often

called overpotentials, η , are defined as the deviation of the cell potential, Ecell, from the equilibrium thermodynamic potential E. The cell potential is obtained by subtracting all overpotentials (losses) from the equilibrium thermodynamic potential as the following expression;

E cell = E − η act − η ohm − η mem − η Diff

(41)

The equilibrium potential is determined using the Nernst equation [27];

( )

( )

1 ⎤ ⎡ E = 1.229 − 0.83 × 10 − 3 (T − 298.15) + 4.3085 × 10 −5 T ⎢ln PH 2 + ln PO2 ⎥ (42) 2 ⎦ ⎣

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The anode and cathode activation overpotentials are calculated from Butler-Volmer equation (33 and 34). The ohmic overpotentials in GDLs and protonic overpotential in membrane are calculated from the potential equations (24) and (40) respectively. The anode and cathode diffusion overpotentials are calculated from the following equations [28];

η Diff ,c =

i RT ⎛ ln⎜⎜1 − c 2 F ⎝ i L,c

⎞ ⎟ ⎟ ⎠

(43)

η Diff ,a =

i RT ⎛ ln⎜⎜1 − a 2 F ⎝ i L,a

⎞ ⎟ ⎟ ⎠

(44)

i L,c =

i L ,a =

2 FDO2 CO2

(45)

δ GDL 2 FD H 2 C H 2

(46)

δ GDL

The diffusivity of oxygen and hydrogen are calculated from the following equations [29];

DO 2

⎛ T ⎞ = 3.2 × 10 ⎜ ⎟ ⎝ 353 ⎠ −5

32

⎛ 101325 ⎞ ⎟ ⎜ ⎝ P ⎠

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DH2

⎛ T ⎞ = 1.1 × 10 ⎜ ⎟ ⎝ 353 ⎠ −4

32

⎛ 101325 ⎞ ⎟ ⎜ ⎝ P ⎠

143

(48)

3. RESULTS AND DISCUSSION

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The governing equations were discretized using a finite volume method and solved using a multi-physics computational fluid dynamics code. Stringent numerical tests were performed to ensure that the solutions were independent of the grid size. A computational quadratic mesh consisting of a total of 4611 nodes and 22851 meshes for airflow-channel design, 3679 nodes and 18030 meshes for planar air-breathing design, and 7185 nodes and 37877 meshes for tubular air-breathing design ware found to provide sufficient spatial resolution (Figure 6). The coupled set of equations was solved iteratively, and the solution was considered to be convergent when the relative error in each field between two consecutive iterations was less than 1.0×10−6. The calculations presented here have all been obtained on a Pentium IV PC (3 GHz, 2GB RAM) using Windows XP operating system.

Figure 6. Computational meshes of PEM fuel cells (quadratic): airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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The values of the electrochemical transport parameters for the base case operating conditions are taken from reference [30] and are listed in Table 1. The geometric and the base case operating conditions are listed in Table 2. It is important to note that because this model accounts for all major transport processes and the modelling domain comprises all the elements of a complete cell, no parameters needed to be adjusted in order to obtain physical results. Table 1. Electrode and membrane properties and parameters. Parameter

Symbol

Value

Electrode porosity

ε kp e

0.4

Electrode thermal conductivity (Ballard AvCarb®-P150)

keff

17.1223

W / m.K

Membrane thermal conductivity

kmem

0.455

W / m.K

Electrode electronic conductivity

λe

180

S /m

Membrane ionic conductivity (humidified Nafion® 117)

λm

17.1223

S /m

Transfer coefficient, anode side

αa

0.5

-

Transfer coefficient, cathode side

αc

1

-

Anode ref. exchange current density

ioref, a

2465.598

A / m2

Cathode ref. exchange current density

ioref, c

1.8081 × 10 3

A / m2

Oxygen concentration parameter

γO

1

-

Hydrogen concentration parameter

γH

.5

-

Entropy change of cathode reaction

ΔS

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Electrode hydraulic permeability

2

2

Unit -

1.7 × 10

−11

-326.36

m2

J / mole.K W / m3

Heat transfer coefficient

β

4 × 10

Fuel cell temperature

Tcell

353.15

K

Protonic diffusion coefficient

DH +

4.5 × 10 −9

m2 / s

Fixed-charge concentration

cf

1200

mole / m3

Fixed-site charge

zf

-1

-

Electro-osmotic drag coefficient

nd

2.5

-

Droplet diameter

D drop

1.0 × 10 −8

m

Condensation constant

C

-

m2

6

Scaling parameter for evaporation

ϖ

1.0 × 10 −5 0.01

Membrane hydraulic permeability

kp m

7.04 × 10 −11

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Table 2. Geometrical and operational parameters for base case operating conditions. Parameter

Symbol

Value

Unit

Channel length Hydrogen flow channel: diameter (tubular design) / height & width (planar design)

L

50 × 10 −3

m

H H2

1 × 10 −3

m

Ambient Air inter section width (planar design)

W Air

1 × 10 −3

m

Land area width (planar design)

Wland

0.5 × 10 −3

m

Electrode thickness (GDL)

δ GDL

0.26 × 10 −3

m

Catalyst layer thickness

δ CL

0.0287 × 10 −3

m

Membrane thickness

δ mem

0.23 × 10 −3

m

Hydrogen reference mole fraction

xHref2

0.84639

-

Oxygen reference mole fraction

xOref2

0.17774

-

Air pressure (Cathode pressure)

Pc

1

atm

Fuel pressure (Anode pressure)

Pa

1

atm

Fuel stoichiometric flow ratio

ξa ψ

2

-

0.79/0.21

-

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Inlet Oxygen/Nitrogen ratio

Polarization curves of the airflow-channel type and ambient air-breathing planar and tubular-shaped PEM fuel cells were compared with the experimental results provided by Wang et al. [31] for the airflow-channel PEM fuel cell in the same operating conditions (Figure 7). Polarization curves of the two models of the air-breathing fuel cells clearly show that higher cell potentials are achieved with the air-breathing tubular design mainly because of lower activation and diffusion potentials. Better gas replenishment at the catalyst sites in tubular design results in lower of activation and diffusion potentials. The tubular-shaped fuel cell does not seem to have higher mass transport losses. The linear behaviour of the voltagecurrent curve for the tubular-shaped fuel cell suggests that the overall overpotential is driven mainly by ohmic losses. In addition, Figure 7 shows that the planar ambient air-breathing fuel cells have a higher cell potential than airflow-channel fuel cells. This is because of the oxygen concentration decreases gradually from the inflow channel to the outflow channel in the airflow-channel fuel cells due to the consumption of oxygen at the catalyst layer. The lower diffusivity of the oxygen along with the low concentration of oxygen in ambient air results in noticeable oxygen depletion near outflow channel areas. Moreover, fresh air coming from the inflow channel has a longer distance to diffuse through to reach these areas. In addition to the polarization curve, the comprehensive three-dimensional CFD model also allows for the assessment of important information about the detail of transport phenomena inside the fuel cell. In order to gain some insight into why the polarization curves are better in the case of the tubular-shaped PEM fuel cell, the multiphase velocity flow field, oxygen and hydrogen distribution, local current densities, temperature distribution, and

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potential distribution are plotted for all designs; (airflow-channel type and planar and tubular air-breathing), for a fixed nominal current density of 0.8 A.cm-2.

Figure 7. Comparison of the models and the experimental polarization curves.

The detailed distribution of oxygen molar fraction is shown in Figure 8. In airflowchannel type, oxygen concentration decreases gradually from the inflow channel to the outflow channel due to the consumption of oxygen at the catalyst layer. In the GDL, oxygen concentration under the land area is smaller than that under the air inlet area. The concentration of oxygen at the catalyst layer is balanced by the oxygen that is being consumed and the amount of oxygen that diffuses towards the catalyst layer driven by the concentration gradient. The lower diffusivity of the oxygen along with the low concentration of oxygen in ambient air results in noticeable oxygen depletion under the land areas. The planar air-breathing design gives more even distribution of the molar oxygen fraction at the catalyst layer than airflow-channel type. At planar air-breathing design, the oxygen mole fraction variation is low enough not to cause diffusive limitations, whereas at airflow-

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channel type the concentration of oxygen under the land areas has already reached near-zero values. The molar oxygen fraction at the catalyst layer increases with more even distribution with planar air-breathing design. This is because of a better gas replenishment at the catalyst sites results in quite uniform distribution for the oxygen to reach the catalyst layer. The tubular air-breathing design gives further even distribution of the molar oxygen fraction at the catalyst layer. At a tubular shape, the oxygen mole fraction variation is very low enough not to cause diffusive limitations. The molar oxygen fraction at the catalyst layer increases with much more even distribution with tubular shape. This is because of a better gas replenishment at the catalyst sites of the tubular shape results in quite uniform distribution for the oxygen to reach the catalyst layer. Due to the relatively low diffusivity of the oxygen compared with that of the hydrogen, the cathode operation conditions usually determine the limiting current density. This is because an increase in current density corresponds to an increase in oxygen consumption.

Figure 8. Oxygen molar fraction distribution: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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The hydrogen molar fraction distribution in the anode side is shown in Figure 9. In general, the hydrogen concentration decreases from inlet to outlet as it is being consumed. However, the decrease is quite small along the channel and the decrease in molar concentration of the hydrogen under the land areas of the planar shapes is smaller than for the oxygen in cathode side due to the higher diffusivity of the hydrogen.

Figure 9. Hydrogen molar fraction distribution: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

Figure 10 shows the local current density (in A.cm-2) distribution at the cathode side catalyst layer. The local current densities have been normalized by divided through the nominal current density in each case (i.e. ic/I). The local current density of the cathode side reaction depends directly on the oxygen concentration. The diffusion of the oxygen towards the catalyst layer is the main impediment for reaching high current densities. Therefore, it can be seen that for tubular air-breathing design, the distribution of the local current density is quite uniform, and for planar air-breathing design, the distribution of the local current density is lightly uniform. This is change for airflow-channel type, where under the land areas a noticeable decrease takes place. It can be seen that for airflow-channel type, a high fraction of the current is generated at the catalyst layer that lies in front of the channel, leading to underutilization of the catalyst under the land areas. This can lead to local hot spots inside the

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membrane electrode assembly. These hot spots can lead to a further drying out of the membrane, thus increasing the electric resistance, which in turn leads to more heat generation and can lead to a failure of the membrane. Thus, it is important to keep the current density relatively even throughout the cell. For optimal fuel cell performance, a uniform current density generation is desirable (as shown in tubular shape), and this could only be achieved in planar design with a non-uniform catalyst distribution, possibly in conjunction with nonhomogeneous gas diffusion layers.

Figure 10. Local current density distribution: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

Thermal management is required to remove the heat produced by the electrochemical reaction in order to prevent drying out of the membrane and excessive thermal stresses that may result in rupture of the membrane. The small temperature differential between the fuel cell stack and the operating environment make thermal management a challenging problem in PEM fuel cells [32]. The temperature distribution inside the fuel cell has important effects on nearly all transport phenomena, and knowledge of the magnitude of temperature increases due to irreversibilities might help preventing failure [32]. Figure 11 shows the distribution of the temperature (in K) inside the cell. The result of all designs shows that the increase in temperature can exceed several degrees Kelvin near the catalyst layer regions, where the electrochemical activity is highest. The temperature peak appears in the cathode catalyst layer, implying that major heat generation takes place in the

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region. In general, the temperature at the cathode side is higher than that at the anode side; this is due to the reversible and irreversible entropy production. The air-breathing design results in more even distribution of the local current density with low fraction than airflowchannel type. Therefore, the maximum temperature gradient appears in the airflow-channel type and the minimum temperature gradient appears in the tubular air-breathing design as can be seen in Figure 11. For an optimum fuel cell performance, and in order to avoid large temperature gradients inside the fuel cell, it is desirable to achieve a uniform current density distribution inside the cell.

Figure 11. Temperature distribution inside the cell: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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Activation overpotentials (in V) distribution for both geometries is shown in Figure 12. The activation overpotential profile correlates with the local current density. For the planar air-breathing design, fresh air coming from the air inlet area has a longer distance to diffuse through to reach the land areas. This fact results in a diminished oxygen concentration at the catalyst sites under the land areas of the planar air-breathing design. Therefore, the planar airbreathing fuel cell leads to a distribution where the maximum is located at the centre of the air inlet area and coincide with the highest reactant concentrations. Better gas replenishment at the catalyst sites in the tubular air-breathing design results in lower activation potentials with quite uniform distribution. In airflow-channel type, the potential distribution at the catalyst layer is a gradient along the flow channel. This is due to the oxygen concentration decreases gradually from the inflow channel to the outflow channel due to the consumption of oxygen at the catalyst layer.

Figure 12. Activation overpotential distribution at the cathode sites: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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To perform a comprehensive comparison study for each components of the cell, two types of ohmic losses that occur in MEA are characterized. These are potential losses due to electron transport through electrodes and potential loss due to proton transport through the membrane. Ohmic overpotential is the loss associated with resistance to electron transport in the GDLs. For a given nominal current density, the magnitude of this overpotential is dependent on the path of the electrons. The potential field (in V) in the cathodic and the anodic gas diffusion electrodes are shown in Figure 13. The potential distributions are normal to the flow inlet, fuel and air, and the side walls. The distributions exhibit gradients due to the non-uniform local current production. For planar shape design, there is a gradient into the land areas where electrons flow into the bipolar plate. For tubular shape design, there is a gradient into the end areas (top of the anode GDL and bottom of the cathode GDL) where electrons flow into the series connections from anode to cathode.

Figure 13. Ohmic overpotential distribution in the anode and cathode GDLs: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

The potential loss in the membrane is due to resistance to proton transport across the membrane from anode catalyst layer to cathode catalyst layer. The distribution pattern of the protonic overpotential is dependent on the path travelled by the protons and the activities in the catalyst layers. Figure 14 shows the potential loss distribution (in V) in the membrane. It can be seen that the potential drop is more uniformly distributed across the membrane in all designs. This is because of the smaller gradient of the hydrogen concentration distribution at the anode catalyst layer due to the higher diffusivity of the hydrogen.

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153

Figure 14. Membrane overpotential distribution across the membrane: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

The variation of the cathode diffusion overpotentials (in V) is shown in Figure 15. Tubular air-breathing design improves the mass transport within the cell and this leads to reducing the mass transport loss. Better gas replenishment at the catalyst sites results in lower and quite uniform distribution of diffusion potentials. For planar air-breathing design, there is a much stronger distribution of diffusion potentials at the catalyst sites, with higher values in front of the air inlet area. This is due to the reduction of the molar oxygen fraction at the catalyst layer under the land areas. In airflow-channel type, the diffusion potential distribution at the catalyst layer is a gradient along the flow channel. This is because of the oxygen concentration decreases gradually from the inflow channel to the outflow channel due to the consumption of oxygen at the catalyst layer. The lower diffusivity of the oxygen along with the low concentration of oxygen in ambient air results in noticeable oxygen depletion near outflow channel areas. Moreover, fresh air coming from the inflow channel has a longer distance to diffuse through to reach these areas.

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Figure 15. Diffusion overpotential distribution at the cathode sites: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

The operation of a fuel cell and the resulting water and heat distributions depend on numerous transport phenomena including charge-transport and multi-component, multiphase flow, and heat transfer in porous media. Multiphase flow is a central issue in PEM fuel cell technology because while water is essential for membrane ionic conductivity, excess liquid water leads to flooding of catalyst layers and gas diffusion layers. Understanding the flow of gas/liquid flows is therefore of major technological as well as scientific interest. The multiphase velocity fields inside the cathodic and anodic gas diffusion layers are shown in Figures 16, 17, 18 and 19 for both gas and liquid phase. The pressure gradient induces bulk gas flow from the hydrogen channel and ambient air into the GDL. While the capillary pressure gradient drives the liquid water out of the gas diffusion layers into the hydrogen flow channel and ambient air. Therefore, the liquid water flux is directed from the GDL into the hydrogen channel and ambient air, i.e., in the opposite direction of the gas-phase velocity, where it can leave the cell. The velocity of the liquid phase, however, is lower than for the gas phase, which is due to the higher viscosity, and the highest liquid water velocity occurs at the corners of the current-

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collector/GDL interface. The liquid water oozes out of the GDL, mainly at the corners of the current-collector/GDL interface.

Figure 16. Gas phase velocity vectors inside the cathode GDL: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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Figure 17. Liquid water velocity vectors inside the cathode GDL: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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  Figure 18. Gas phase velocity vectors inside the anode GDL: airflow-channel design (upper); planar airbreathing design (middle); tubular air-breathing design (lower).

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  Figure 19. Liquid water velocity vectors inside the anode GDL: airflow-channel design (upper); planar air-breathing design (middle); tubular air-breathing design (lower).

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4. CONCLUSIONS The increasing number of portable electronic devices on the market today such as laptop computers, cellular phones, PDAs, remote-controlled toys, is creating a demand for improved, more environmentally friendly technologies in energy storage and conversion. Air-breathing PEM fuel cells have the potential to supplant batteries in portable electronic devices because of increased cost tolerances for fuel cells at this scale, their rapid recharges, and their higher energy densities. The development of physically representative models that allow reliable simulation of the processes under realistic conditions is essential to the development and optimization of fuel cells, the introduction of cheaper materials and fabrication techniques, and the design and development of novel architectures. The difficult experimental environment of fuel cell systems has stimulated efforts to develop models that could simulate and predict multi-dimensional coupled transport of reactants, heat and charged species using computational fluid dynamic (CFD) methods. The strength of the CFD numerical approach is in providing detailed insight into the various transport mechanisms and their interaction, and in the possibility of performing parameters sensitivity analyses. The results of CFD analyses are relevant in: conceptual studies of new designs, detailed product development, troubleshooting, and redesign. CFD analysis complements testing and experimentation, by reduces the total effort required in the experiment design and data acquisition. Full three-dimensional computational fluid dynamics models of planar and novel tubularshaped air-breathing PEM fuel cell have been developed and presented in detail in this Chapter. In addition, the typical airflow-channel PEM fuel cell has been modelled to use as a basis for comparison. These models provides valuable information about the transport phenomena inside the air-breathing PEM fuel cell such as reactant gas concentration distribution, multiphase velocity flow field, temperature distribution, potential distribution in the membrane and gas diffusion layers, activation overpotential distribution, diffusion overpotential distribution, and local current density distribution. The results show that higher power densities are achieved with the tubular design mainly because of lower activation and diffusion overpotentials. Better gas replenishment at the catalyst sites in tubular design results in lower and quite uniform distribution of activation and diffusion potentials. A tubular shape evens out the local current density distribution. For a planar shape a much higher fraction of the total current is generated under the air inlet area. Therefore, the maximum temperature gradient appears in the planar shape design. The models are shown to be able to: [1] [2] [3]

understand the many interacting, complex electrochemical and transport phenomena that cannot be studied experimentally; identify limiting steps and components; and provide a computer-aided tool for the design and optimization of future fuel cells to improve their lifetime with a much higher power density and lower cost.

They can provide a solid basis for optimizing the geometry of the PEM fuel cell stack running with a passive mode.

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ACKNOWLEDGMENTS My great appreciation is expressed to International Technological University (ITU), London, UK for providing available facilities. My gratitude and appreciation is due to my wife for her patience, care, and support during the period of preparing this Chapter. Most of all, I want to thank my parents for their unconditional support throughout all the years of my education. This study would not have been possible without them. Many thanks also are due to my brother and sister. Maher A.R. Sadiq AL-Baghdadi

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Nomenclature Symbol

Description

Units

a

Water activity

-

AMEA

Area of the MEA

m2

Ach

Cross sectional area of flow channel

m2

cW

Water content

-

C

Condensation constant

-

Cf

Fixed charge concentration

mol.m −3

CH2

Local hydrogen concentration

mol.m −3

C Href2

Reference hydrogen concentration

mol.m −3

C O2

Local oxygen concentration

mol.m −3

C Oref2

Reference oxygen concentration

mol.m −3

Cp

Specific heat capacity

J .kg −1 .K −1

D

Diffusion coefficients

m 2 .s −1

DH +

Protonic diffusion coefficient

D drop

Diameter of droplet water

m 2 .s −1 m

E

Reversible cell potential

volts

E cell

Cell operating potential

volts

E fc

Thermodynamic efficiency of the cell

-

F

Faraday's constant

G

Gibb's free energy

96487C.mol −1 J

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Nomenclature (continued)

g

Specific Gibb’s free energy

J .mol −1

h

Specific enthalpy

J .mol −1

I

Cell operating (nominal) current density

A.m −2

ia

Anode local current density

A.m −2

ic

Cathode local current density

A.m −2

ioref,c

Anode reference exchange current density

A.m −2

ioref,a

Cathode reference exchange current density

A.m −2

i L,a

Anode local limiting current density

A.m −2

i L,c

Cathode local limiting current density

A.m −2

k

Gases thermal conductivity

W .m −1 .K −1

keff

Effective electrode thermal conductivity

W .m −1 .K −1

k gr

Graphite thermal conductivity

W .m −1 .K −1

kmem

Membrane thermal conductivity

W .m −1 .K −1

k xm

Mass transfer coefficient

mol.m −2 .s −1

Kp

Hydraulic permeability

m2

LHV H 2

Lower heating value of hydrogen

J .kg −1

m& phase

Mass transfer in the form of: evaporation

kg.s −1

M

Molecular weight of mixture gases

kg.mol −1

M H2

Molecular weight of hydrogen

kg.mol −1

M H 2O

Molecular weight of water

kg.mol −1

M O2

Molecular weight of oxygen

kg.mol −1

NW

Net water flux across the membrane

kg.m −2 .s −1

nd

Electro-osmotic drag coefficient

-

ne P

Number of electrons transfer

-

Pressure

Pa

Pc

Capillary pressure

Pa

Psat q&

Water saturation pressure

Pa

Heat generation

W .m −2

m& phase = m& evap and condensation m& phase = m& cond

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Maher A.R. Sadiq Al-Baghdadi Nomenclature (continued)

rg

Volume fraction of the gas phase

-

rl R

Volume fraction of the liquid phase

-

Universal gas constant

8.314 J .mol −1 .K −1

s

Specific entropy

J .mol −1 .K −1

sat

Saturation

-

T u

Temperature

K

Velocity vector

m.s −1

v∞

Free-stream velocity

m.s −1

Wcell

Cell power density

W .m −2

xi

Molar fraction

-

yi

Mass fraction

-

zf

Fixed-site charge

-

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Greek Symbols Description Symbol

Description

Units

αa

Charge transfer coefficient, anode side

-

αc β τ

Charge transfer coefficient, cathode side

-

Modified heat transfer coefficient Chemical potential

W .m −3 .K −1 J .mol −1

Electrochemical potential

J .mol −1

Hydrogen concentration parameter

-

Oxygen concentration parameter

-

ΔH evap

Enthalpy of evaporation

J .kg −1

ΔS

Entropy change of cathode side reaction

δ CL

Catalyst layer thickness

J mole −1 K −1 m

δ GDL

Gas diffusion layer thickness

m

δ mem ε

Membrane thickness

m

Porosity

-

τ γH

γO

2

2

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CFD Models for Analysis and Design of Ambient Air-Breathing PEM Fuel Cells Greek Symbols Description (continued)

ξ η

Stoichiometric flow ratio

-

Overpotential

volts

λe

Electrode electronic conductivity

S .m −1

λm μ

Membrane ionic conductivity

S .m −1

Viscosity

ρ

kg.m −1 .s −1

Density

φGDL

Electric potential inside the gas diffusion layer

kg.m −3 volts

φ mem

Electric potential inside the membrane

volts

ℜ

Relative humidity of inlet fuel and air

%

ψ

Inlet Oxygen/Nitrogen ratio

-

ϕ

Local relative humidity of the gas phase

-

ϖ

Scaling parameter for evaporation

-

σ

Surface Tension

N .m −1

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Abbreviations Definition CFD CL GDL MEA MEM PEM sat

Computational Fluid Dynamics Catalyst Layer Gas Diffusion Layer Membrane Electrode Assembly Membrane Proton Exchange Membrane Saturation

Subscript Definition

a c g l w

Anode Cathode Gas phase Liquid phase Water

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REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14]

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[15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25] [26]

O’Hayre, R; Fabian, T; Litster, S; Prinz, F.B.; Santiago, J.G. Journal of Power Sources. 2007, 167(1), 118–129. Litster, S.; Pharoah, J.G.; McLean, G.; Djilali, N. Journal of Power Sources. 2006, 156(2), 334–344. Hwang, J.J.; Wu, S.D.; Pen, R.G.; Chen, P.Y.; Chao, C.H. Journal of Power Sources. 2006, 160(1), 18–26. Ying, W.; Sohn, Y.; Lee, W.; Ke, J.; Kim, C. Journal of Power Sources. 2005, 145(2), 563–571. Wang, Y.; Ouyang, M. Journal of Power Sources. 2007, 164(2), 721–729. Litster, S.; Djilali, N. Electrochimica Acta. 2007, 52(11), 3849–3862. Rajani, B.P.M.; Kolar, A.Kumar. Journal of Power Sources. 2007, 164(1), 210–221. Zhang, Y.; Pitchumani, R. Int. Journal of Heat & Mass Transfer. 2007, 50(23-24), 4698–4712. Zhang, Y.; Mawardi, A.; Pitchumani, R. Journal of Power Sources. 2007, 173(1), 264– 276. Paquin, M.; Frechette, L. Journal of Power Sources. 2008, 180(1), 440–451. Noponen, M.; Mennola, T.; Mikkola, M.; Hottinen, T.; Lund P. Journal of Power Sources. 2002, 106(1-2), 304–312. Schmitz, A.; Tranitz, M.; Wagner, S.; Hahn, R.; Hebling, C. Journal of Power Sources. 2003, 118(1-2), 162–171. Hottinen, T.; Mikkola, M.; Lund, P. Journal of Power Sources. 2004, 129(1), 68–72. Jeong, S.; Cho, E; Kim, H; Lim, T; Oh, I; Kim, S. Journal of Power Sources. 2006, 159(2), 1089–1094. Fabian, T.; Posner, J.; O’Hayre, R.; Chac, S.; Eaton, J. Prinz, F.; Santiago, J. Journal of Power Sources. 2006, 161(1), 168–182. Rosa, D.; Pinto, D.; Silva, V. Silva, R.; Rangel, C. Int. Journal of Hydrogen Energy. 2007, 32(17), 4350–4357. Jung, U.; Jeong, S., Park, K.; Lee, H.; Chun, K.; Choi, D.; Kim, S. Int. Journal of Hydrogen Energy. 2007, 32(17), 4459–4465. Fuller, E.N.; Schettler, P.D.; Giddings, J.C. Ind. Eng. Chem. 1966, 58(5), 18-27. Berning, T.; Djilali, N. Journal of Electrochem. Soc. 2003, 150(12), A1589-A1598. Coker, A.K. FORTRAN programs for chemical process design, analysis, and simulation; ISBN: 0-88415-280-4; Gulf Publishing Company: Houston, Texas, 1995. De La Rue, R.E.; Tobias, C.W. Journal of Electrochem. Soc. 1959, 106(9), 827–836. Al-Baghdadi, M.; Al-Janabi, H. Energy Conversion and Management. 2007, 48(12), 3102-3119. Sui, P.; Djilali N. Journal of Fuel Cell Sci. Tech. ASME. 2005, 2(1), 149-155. Siegel, N.P.; Ellis, M.W.; Nelson, D.J.; von Spakovsky, M.R. Journal of Power Sources. 2004, 128(2), 173–184. Hu, M.; Gu, A.; Wang, M.; Zhu, X.; Yu, L. Energy Conversion and Management. 2004, 45(11-12), 1861–1882. Hu, M.; Gu, A.; Wang, M.; Zhu, X.; Yu, L. Energy Conversion and Management. 2004, 45(11-12), 1883–1916.

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[27] Al-Baghdadi, M. Renewable Energy Journal. 2008, 33(6), 1334-1345. [28] Hirschenhofer, J.H.; Stauffer, D.B.; Engleman, R.R.; Klett; M.G. Fuel Cell Handbook; DOE/NETL-2002/1179 (DE-AM26-99FT40575); U.S. Department of Fossil Energy, Morgantown Energy Technology Center, Morgantown WV, 2002. [29] Um, S; Wang, C.Y. Journal of Power Sources. 2004, 125(1), 40–51. [30] Al-Baghdadi, M. Fuel Cell Research Trends Textbook, Chapter 7: PEM Fuel Cell Modeling. NOVA SCIENCE PUBLISHERS, INC, 400 Oser Avenue, Suite 1600, Hauppauge, NY 11788 USA. 2007, pp. 273-379. ISBN: 978-1600216695. [31] Wang, L.; Husar, A.; Zhou, T.; Liu, H. Int. Journal of Hydrogen Energy. 2003, 28(11), 1263– 1272. [32] Al-Baghdadi, M. Fuel Cell. 2008, 8(1), 34-39.

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In: Polymer Electrolyte Membrane Fuel Cells… Editors: R. Esposito

ISBN: 978-1-60692-773-1 ©2009 Nova Science Publishers, Inc.

Chapter 5

NEW TRENDS IN THE DEVELOPMENT OF PEMFC CATALYSTS Alevtina L. Smirnova* University of Connecticut, Connecticut Global Fuel Cell Center, Storrs, CT- USA

ABSTRACT

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To decrease the cost of the PEMFC, the properties of the three major components of the catalyst layers should be simultaneously considered including the electrocatalyst, the catalyst support and the polymer electrolyte. The noble group metal electro-catalysts are extensively used for PEMFCs however, high cost of these materials significantly restricts their commercial application. This chapter describes three major groups of PEMFC catalysts including Platinum Group Metal (PGM), platinum alloy and non-noble metal catalysts. The catalytic, chemical, and electrochemical properties of the catalysts are discussed regarding their catalytic activity, electrochemical performance, and stability in oxidizing PEMFC cathode environment. Novel cost-effective and innovative methods for synthesis of these catalysts are reviewed and analyzed in terms of catalyst performance optimization. Specific emphasis is made on carbon and inorganic alternative supports for PEMFC catalysts that should combine chemical stability with high electronic conductivity. Among carbon supports carbon nanotubes, high surface area carbons, carbon blacks and carbon aerogels are described. In contrast to other carbon supports carbon aerogels synthesized in supercritical (SC) conditions have certain advantages including controllable porosity, absence of impurities, and narrow pore size. In this chapter the recent progress achieved in nano-structured carbon aerogels, such as ResorcinolFormaldehyde (RF) or Polyimide (PI) aerogels is reviewed in terms of chemical composition of the precursor gels, xerogel-aerogel structure, synthesis conditions, properties, and morphology. Concentration of micro- and meso-pores, and thus, configuration and texture of carbon aerogels, is reviewed depending on the structure of metallo-organic precursor, SC environment, and carbonization conditions. As an important component of the PEMFC catalysts layer, low molecular weight polymer electrolytes and polymer electrolytes modified with heteropolyacids (HPAs) *

Home:5 Michalec Rd., Stafford Springs CT 06076-4505 USA Work: University of Connecticut/Connecticut Global Fuel Cell Center,44 Weaver Rd, Unit 5233, Storrs, CT 06269 USA

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Alevtina L. Smirnova with improved catalytic activity and conductivity are discussed considering their influence on the overall catalytic activity of the PEMFC cathode. It can be concluded that synergistic combination of the novel materials, such as nano-structured catalysts alloys or core-shell structures, modified carbon or inorganic supports combined with improved polymer electrolytes should significantly increase the catalytic activity and overall PEMFC performance moving PEMFC toward large scale commercialization.

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1.1 INTRODUCTION The fossil fuel depletion and environmental pollution are the major concerns that initiate the alternative clean energy solutions. Among them fuel cells, well known as energy conversion and power generation devices are considered as the most promising alternative. Among different types of fuel cells Proton Exchange Membrane Fuel Cells also known as Polymer Electrolyte Membrane Fuel Cells (PEMFCs) are the most developed in both fundamental and application areas [1] such as commercial, portable, residential, automotive and in distributed power generation [2]. PEMFCs also gained importance as energy accumulators for photovoltaic and wind power sources [3]. In comparison to other fuel cell types, PEMFCs operating in the temperature range of 20140oC have limited capacity for heat and power applications. However in comparison to SOFC they have certain advantages, such as simplicity of operation, zero emissions and potential for low capital and maintenance cost. PEMFCs may be operated at high current densities resulting in fast startup, compact and light weight design, and absence of corrosive fluids since the only liquid present in the cell is water. PEMFCs are capable of operation in the broad range of pressures starting from ambient and reaching 3000psi. A disadvantage associated with PEMFCs that prevent them from large-scale manufacturing is high cost of Pt catalysts that are required as promoters for the electrochemical reaction. PEMFCs consist of a polymer electrolyte membrane and two electrodes, both of which are comprised of expensive carbon supported Pt-Group Metals (PGMs) or metal alloys. As a result of the electrochemical reaction on the anode, the hydrogen molecules split into hydrogen cations (protons) and electrons. Through the external circuit the electrons are transported to the cathode. Protons are transferred through the polymer membrane and on the cathode side combine with oxygen forming electrons and water and releasing heat. Slow oxygen reduction reaction (ORR) kinetics on the cathode is the major concern since it requires about 0.3-0.4mg/cm2 of PGM catalysts which translates into a high PEMFC stack cost. Due to the low operation temperature, PEMFCs require purified H2 to prevent poisoning of PGM catalysts from harmful air and fuel impurities, such as carbon monoxide, sulfur, ammonia and different anions and cations. In general, PEMFC can operate on hydrogen-rich fuel or reformate generated from coal, natural gas, gasoline, landfill gas, or alcohol, but it is a time consuming and expensive fuel processing. The poisoning effect of some of the impurities, such as CO, depends on temperature. For example PEMFC operating on H2 below 80oC can tolerate only a few parts per million (ppm) of CO. However, at higher temperatures (>140oC) CO is less harmful. If a hydrocarbon, for instance a natural gas is used as a fuel, the reformation according to the reaction CH4 + H2O =

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ls eta e m films as d b thin an nd GM s a f P ture n o uc itio str os ell ep -sh 9D ore 9C

9C 9S arb SC urf on g 9T NH ace rap nit hit 9 hin 3 rid iza lay ati tio on n er ,e me .g. tal wi nit th rid ec oa tin gs

3H2 + CO is necessary followed by the gas-shift reaction where CO + H2O = H2 + CO2. Any sulfur compounds present in a hydrocarbon fuel have to be removed to a concentration of 3000m2/g), controllable nano-scale pore structure, narrow pore-size distribution and high electrical conductivity (25-100 S/cm). Their remarkable mechanical and thermal properties [35, 36] are derived from a three-dimensional interconnected nano-network structure that provides both molecular accessibility and rapid mass diffusion. Due to the optimization of the synthesis procedures the cost of CAs reaches

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Alevtina L. Smirnova

now $3 per square foot as insulation material [37,38] or 9 cents per 1000L for water purification [39], which is comparable with the cost of high surface area carbons. Besides, this cost can be reduced further by using inexpensive byproducts of oil-shell industry [40]. Carbon aerogels are derived from organic aerogels originally discovered by S. Kistler in 1931 [41] and later studied by Pekala [42,43]. They are represented by open-cell cross-linked polymers with interconnected pores created through a sol-gel polymerization process. The four major steps (Figure 3) are considered in CA synthesis such as sol-gel formation, solvent exchange, supercritical or ambient pressure drying and pyrolysis, all of which significantly affect the total surface area, microstructure, and porosity of the material. In the first step (Figure 4) an organic aerogel is produced by drying the sol-gel so that the fragile cross-linked internal structure does not collapse thus giving aerogel the highest known surface area per unit volume.

1

2

Gelation

Solvent Exchange

Catalyst

H20 acetone

4

3

Drying

Pyrolysis

scCO2

N2/H2 o 700-900 C

Figure 3. Four major steps in aerogel synthesis: gelation/crosslinking, solvent exchange, drying, and sintering. O

O

C

C

C

C

O

O O

O HO C HN C

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O polyamic acid

H2 N

NH2

O O C OH C NH O

H2 O AA/PY

n Chemical Imidization

N

O

O

C

C

C O

C O

N

polyimide

     

Figure 4. The first step of cross-linking in aerogel synthesis for: a) Resorcinol-formaldehyde (RF) and b) Polyimide (PI) aerogel.

The amount and type of the catalyst, such as sodium carbonate or acetic acid used in the polymerization reaction defines the size, shape and connectivity of the primary network particles. The latter can be reached by using low catalyst (Na2CO3) concentrations in the polymerization reaction. Alternatively, the acid can be used to generate microscale porous structures. In the polymerization process (Figure 5) the formaldehyde and resorcinol components can be changed in the range of 0.4Au(Cu)>Au(Co)>A(Fe) and Pt>Pt(Fe)≥Pt(Cu)≥Pt(Ni)>Pt(Co), which is in correlation with the corresponding values of the energies of the “d-band canter” rather than degree of the d-orbital filling. Sn-shell and Pt-core structures were produced [88] on highly oriented pyrolytic graphite (HOPG). The formation of various PtSn bimetallic structures was attributed to the different surface energy of Pt and Sn, and the interaction of the metals with the HOPG surfaces. Supported on multi-wall carbon nanotubes (MWCNTs) Co and Pt-Ru catalysts [89] with 3.7nm Pt/Ru-shell and about 20nm in diameter Co-core were synthesized by substitution reaction between Co and Pt-Ru alloy possess high electrochemical activity for methanol oxidation that was attributed to high surface of Pt-Ru and their good dispersion on Co particles. Another core–shell structure consisting of a Pt-rich core and Ru-rich shell supported on carbon (PtRu/C) for direct methanol fuel cells was synthesized by a polyol reduction method [90]. In comparison to commercial catalysts the obtained activity was six times as much that was explained by the correlation between the catalytic activity and the Pt–Ru atomic pair frequency occurring on the particle surface, which supports the “bi-functional mechanism”.

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A different example of core –shell structure was reported recently that comprised spherical carbon particles with macroporous core (260 nm) and mesoporous shell (40 nm) [91]. They were used as an anode catalyst support in direct formic acid fuel cell (DFAFC) for Pt50Ru50 (60 wt.%) catalysts and exhibited almost two times the power density in comparison to the commonly used catalyst support carbon black Vulcan XC-72.

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1.3.2. Non-Noble Metal Catalysts In recent years many attempts were undertaken to develop non-noble metal catalysts [92,93]. Currently, a few approaches for synthesis of non-noble metal catalysts are active, among them intercalation compounds, catalysts based on nitrogen enriched carbon supports, and catalysts based on macrocyclic compounds. Carbon supports intercalated with different transition metals, such as Co, Cu, Fe were synthesized by reducing the corresponding metal precursors. Various alkali metals [94] were intercalated into the structure of single-wall carbon nano-tubes and pyrolytic graphites, however, these metals did not possess noticeable catalytic activity. Cu-intercalated carbons were reported by Shioyama et al. [95] that did not show any decrease in catalytic activity due to dissolution or corrosion in acidic media. Modification of carbon support with poly-pyrrole as a chelating polymer possessing ability to form complexes with transition metal cations [96] resulted in high catalytic activity of cobalt catalyst and stable performance as fuel cell cathode catalyst reaching the turnover rate of 0.83s-1 and power density in H2/O2 with 2atm backpressure of 150 mW/cm2 at 0.4 V. In other approaches a source of nitrogen in the form of ammonia or nitrogen compound for modification of carbon support with further intercalation by transition metals was suggested [97] followed by heat treatment in the range of 600-800oC enhancing both catalytic activity and stability. The existing phenomenon was explained [98] by presence of catalytic sites, e.g. FeN4/C and FeN2/C. In comparison to heat treated catalysts, non-heat treated Ni2+ and Pd2+ α-diimine ligands [99] indicated turnover frequencies in ethylene polymerization up to 4×106 mol approaching the activity of metallocenes. Another group of transition metal dihalides, e.g. bis(organilimino) pyridine FeI2+ and Co2+ precursors also known and their activity can be changed by modification of the ligand structure and coordination geometry at the iron center depending on substitution pattern of the N-aryl groups. In addition high catalytic activity of non-noble metal HyPerMecTM DMFC catalysts prepared by reduction of transition metal cations chemically attached to synthetic resins and supported on Vulkan XC-72 was demonstrated. The properties of the non-noble metal catalysts can be also improved by chemical modification of the carbon active groups on the surface of carbon, such as carboxyl, quinone, hydroxyl, and carbonyl using traditional basic (NH4OH) or acidic treatment (HNO3) or scNH3 followed by deposition and chemical stabilization of the transition metal porphyrin-like macro-cycles that posses ORR catalytic activity. During the last decade significant progress in the development of non-noble catalysts such as heat-treated Fe– and Co–N/C catalysts was made towards catalyst synthesis methodologies and understanding of the electro-oxidation mechanism. It was confirmed that

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the enhanced performance of the catalysts strongly depends on the carbon support, the source of metal and nitrogen, and the thermal treatment conditions [100]. A heat-treatment step was identified as important for catalytic activity; however stability improvements are still necessary for metal-based catalysts in acidic PEMFC environment. Regarding synthesis of non-heat treated stabilized catalysts; many studies were performed to develop non-PGM cathode catalysts. The first non-PGM catalysts based on cobalt phthalocyanine organometallic complexes and chalcogenides were discovered by Jasinski in 1964 [101]. Since then many organometallic complexes were investigated but none of them can be considered as a Pt alternative. As non-noble metal catalysts a large variety of inorganic oxides, nitrides, and carbides was investigated among them tungsten carbide with tantalum addition, tantalum oxynitride, zirconium oxide, titanium oxide, zirconium oxynitride, chromium and tantalum nitrides [102], some of them demonstrating certain stability in acidic solutions [103]. A number of tantalum (oxy)nitrides (TaOxNy) was also investigated as new cathode materials for PEMFCs without platinum. One of the investigated Ta oxynitrides (Ta3N5) demonstrated quite low ionization potential and catalytic activity for the ORR [104] though much smaller in comparison to PGMs.

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1.3.3. Supercritical Fluids For Synthesis And Surface Modification Supercritical fluids allow synthesis of different types of nano-particles since the chemical and physical properties of the solvent can be varied with temperature or pressure, both of which can affect the degree of supersaturation and nucleation [105]. Carbon dioxide is known as the most frequently used solvent for supercritical extraction and chemical reactions, however ammonia has been used by far less [106]. In comparison to inert scCO2, scNH3 has strong reducing ability and can be used for surface modification or as a reduction media for deposition of various organometallic and macrocyclic compounds. Ammonothermal synthesis is known to be widely used for nitridation and surface modification, for example for nitridation of vanadium aerogels [107] or synthesis of Ni3N in ScNH3 from Ni(NH3)6Cl2 in presence of NaNH2 at 250oC and 200MPa [108]. Synthesis of nano-structured materials such as Cr-, Co-, Fe-, Cu-, and Ni-nitrides [109 ] was performed in scNH3-methaqol mixtures based on thermal decomposition of metal precursors in a temperature range of 170-290oC at about 16MPa.

1.4. ELECTROLYTE MATERIALS FOR CATALYST MODIFICATION 1.4.1. Low Molecular Weight Electrolyte In PEMFC catalyst layers polymer electrolytes -sulfonated tetrafluoroethylenes are used to increase ionic conductivity and the corresponding catalyst utilization, since only the area of the metal catalyst in contact with electrolyte is active. The proton conductivity of polymer electrolytes (Figure 10) is favored by hydration and strongly depends on the history of the membrane, the operating temperature, and the

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electrolyte environment. The values of n, x, and m can be varied to produce materials of different equivalent weight EW, where EW is the number of grams of the polymer per mole of sulfonate sites. The hydrophobic fluorocarbon chains and the hydrophilic sulfonic groups are arranged to maximize the interaction between the similar fragments. The polymer also contains the water of hydration and protons that are produced on the anode. This results in formation of inverted micelles or ion clusters containing the hydrated ionic phase, which are embedded in the fluorocarbon phase. It is assumed that proton transport occurs between the clusters by proton movement between the fixed sulfonic groups [110].

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Figure 10. Chemical structure of tetrafluoroethylene Nafion®

The catalytic activity and long term stability of PEMFC cathode can be further improved by using the stabilized short-side-chain copolymers (EW of 820 and 790 g/eq SO3H) of tetrafluoroethylene and sulfonyl fluoride vinyl ether [111]. The advantages of low EW short side chain polymer in comparison to Nafion®, such as better proton conductivity [112], higher crystallinity, and higher chemical [113] and mechanical stability [114, 115, 116] improve the overall PEMFC cathode conductivity and performance [117, 118]. As demonstrated earlier [119] the effect of adding low equivalent weight Nafion is more pronounced at lower relative humidities (45%RH) corresponding to higher temperatures (120oC). However, this effect is also important for PEMFCs operating in dry conditions at lower temperatures (70-80oC) which is essential for automotive applications that can be optimized experimentally by analyzing catalytic activity vs. mass transport. It was demonstrated that in presence of transition metal cation impurities or Cl- anions the so-called “unzipping degradation reaction” [120] starts from the carboxylic end groups and disaggregates the main chain of the polymer. In this reaction fluorine and H2O2 can be produced which is harmful for both the electrolyte and the catalyst. Significant (1/4) decrease in fluorine release observed for stabilized short-side-chain copolymer, e.g. Aquivion® (Figure 11) in comparison to Nafion®1100 can significantly improve stability of the cathode due to reduced number of carboxylic groups (-COOH) and decreased H2O2 production. Finally, the advantage of polymer with higher glass transition temperature in comparison to Nafion ensures better adhesion of the catalyst layer during the MEA manufacturing.

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(CF2CF)n- (CF2CF2)mO-(CF2CFx-O)j-(CF2)k-SO3H Figure 11. Chemical structure of the perfluorinated low equivalent weight (EW 830) Aquivion®.

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1.4.2 Catalyst Layers Modified With Heteropolyacids Heteropolyacids (HPAs) are known to possess exceptionally high conductivities even at elevated temperatures due to the special crystal structure (Figure 12) containing incorporated water molecules. It is known that their conductivity in solid state is comparable to the conductivity of sulfuric acid. Furthermore, these compounds maintain high conductivity even at temperatures exceeding boiling point of water. To increase the conductivity of the catalyst layer at lower RH humidities the additional modification of the sulfonated tetrafluoroethylene polymer electrolytes with heteropolyacid (HPAs) [121] can be provided, thus increasing water content, improving cathode performance and impeding degradation effects at start-up and shut-down conditions. Since the kinetics of the oxygen reduction reaction (ORR) start to decrease below 60-70%RH due to low proton activity [122], incorporation of heteropolyacids [123, 124, 125, 126] in the form of insoluble salts is essential for further improvements of the polymer electrolyte, e.g. Hyflon® or Nafion ® at low RH. The effect of HPA, such as silicotungstic (STA) and phosphotungstic (PTA) [127] acid in the form of insoluble salts was tested and analyzed in terms of crystal structure of the heteropolyacids, ionic conductivity, and solubility in water produced on the cathode. The improvement of the cathode performance was explained by the structure of heteropolyacid insoluble salts that are capable to preserve water at low relative humidities and posses conductivities close to the conductivity of sulfuric acid in a dry state [128, 129]. Furthermore, since some of the HPAs, such as molybdenum based HPAs (H5PMo10V2O40) are active in the ORR additional modification of the catalyst layer with insoluble HPA salts can improve catalytic activity. To decrease STA solubility in water, Cs-STA insoluble salt was incorporated into the catalysts layer structure. In terms of Cs-HPA stability an additional approach of using an inorganic matrix was considered [130] using the synthesized inorganic nitrides and carbides. Among 17 investigated V, Fe, and Mn substituted heteropolyacids only molybdenum based acids were active for the ORR substituted as non-Pt ORR catalysts [131]. In the fuel cell environment no HPA reduction activity was recorded above 0.55V that indicated that the catalyst must first be reduced in situ by 4e− before the HPA can reduce oxygen The potential at which the HPA can be reduced has been determined as the limiting factor. The power densities for PEMFC operating on H2 and O2 of 67mW/cm2 and 86mW/cm2 at 0.2V were obtained by using adsorbed and covalently bonded to the carbon surface H5PMo10V2O40 correspondingly.

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Figure 12. The investigated structures of the HPAs: (a) Keggin, (b) Wells–Dawson and (c) Wells– Dawson Sandwich (With permission of the Electrochimica Acta, Stanis 2008).

1.5. CONCLUSION

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An interest in the development of PEMFC catalysts is shifting now toward nanostructured PGM and non-noble metal catalysts supported on various carbon and inorganic supports. The role of the support and the structure of metal nanoparticles, their interaction, stability, and chemical compatibility with electrolyte material are still not completely understood. Thus, more efforts should be undertaken to achieve the goals in synthesis of the cathode catalysts with high catalytic activity reaching our goals in fuel cell efficiency and power generation.

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[74] O. Bourbia, S. Achour, N. Tabet, M. Parlinska, A. Harabi Effect of tantalum addition on microstructure and optical properties of TiN thin films, Thin Solid Films, 515(17) (2007) 6758-6764. [75] J. Bonitz, S. E. Schulz, and T. Gessner, CVD TiN layers as diffusion barrier films on porous SiO2 aerogel, Microelectronic Engineering, 70(2-4) (2003) 330-336. [76] C. Folgar, D. Folz, C. Suchicital, D. Clark, Microstructural evolution in silica aerogel, Journal of Non-Crystalline Solids, 353(16-17) (2007) 1483-1490. [77] E. P. Ambrosio, C. Francia, M. Manzoli, N. Penazzi, P. Spinelli, Platinum catalyst supported on mesoporous carbon for PEMFC, International Journal of Hydrogen Energy, 33(12) (2008) 3142-3145. [78] A. F. Gullá, L. Gancs, R. J. Allen, S. Mukerjee, Carbon-supported low-loading rhodium sulfide electrocatalysts for oxygen depolarized cathode applications, Applied Catalysis A: General, 326( 2) (2007) 227-235. [79] M. Chisaka, H. Daiguji, Design of ordered-catalyst layers for polymer electrolyte membrane fuel cell cathodes, Electrochemistry Communications, 8(8) ( 2006) 13041308.i [80] Y. Zhang, D. Kang, C. Saquing, M. Aindow, and C. Erkey, Supported Platinum Nanoparticles by Supercritical Deposition, Ind. Eng. Chem. Res, (2005) 44, 4161 – 4164. [81] J. Zhou, J. He, Y. Ji, G. Zhao, C. Zhang, X. Chen, T. Wang, Influence of Hierarchical Porosity in Carbon Material on Electrocatalytic Property of Supported Pt Nanoparticles, Acta Physico-Chimica Sinica, 24(5) (2008) 839-843. [82] A. Seo, J. Lee, K. Han, H. Kim, Performance and stability of Pt-based ternary alloy catalysts for PEMFC, Electrochimica Acta, 52( 4) ( 2006) 1603-1611. [83] E. Antolini, J. R.C. Salgado, E. R. Gonzalez, The stability of Pt–M (M=first row transition metal) alloy catalysts and its effect on the activity in low temperature fuel cells: A literature review and tests on a Pt–Co catalyst, Journal of Power Sources, 160( 2) (2006) 957-968. [84] R.K. Raman, A.K. Shukla, A. Gayen, M.S. Hegde, K.R. Priolkar, P.R. Sarode, S. Emura, Tailoring a Pt–Ru catalyst for enhanced methanol electro-oxidation, Journal of Power Sources, 157(1) (2006) 45-55. [85] Z.D. Wei, Y.C. Feng, L. Li, M.J. Liao, Y. Fu, C.X. Sun, Z.G. Shao, P.K. Shen, Electrochemically synthesized Cu/Pt core-shell catalysts on a porous carbon electrode for polymer electrolyte membrane fuel cells, Journal of Power Sources, 180( 1) (2008) 84-91. [86] I. Morjan, I. Soare, R. Alexandrescu, L. Gavrila-Florescu, R.-E. Morjan, G. Prodan, C. Fleaca, I. Sandu, I. Voicu, F. Dumitrache, E. Popovici, Carbon nanotubes grown by catalytic CO2 laser-induced chemical vapor deposition on core-shell Fe/C composite nanoparticles, Infrared Physics & Technology, 51( 3) (2008) 186-197. [87] S. Papadimitriou, A. Tegou, E. Pavlidou, S. Armyanov, E. Valova, G. Kokkinidis, S. Sotiropoulos, Preparation and characterisation of platinum- and gold-coated copper, iron, cobalt and nickel deposits on glassy carbon substrates, Electrochimica Acta, 53( 22) (2008) 6559-6567.

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[88] Y. Yao, Q. Fu, Z. Zhang, H. Zhang, T. Ma, D. Tan, X. Bao, Structure control of Pt–Sn bimetallic catalysts supported on highly oriented pyrolytic graphite (HOPG), Applied Surface Science, 254 (13) (2008) 3808-3812. [89] H. Zhao, L. Li, J. Yang, Y. Zhang, Co and Pt–Ru core-shell nanoparticles supported on multiwalled carbon nanotube for methanol oxidation, Electrochemistry Communications, In Press, Available online (2008). [90] H. Nitani, T. Nakagawa, H. Daimon, Y. Kurobe, T. Ono, Y. Honda, A. Koizumi, S. Seino, T. A. Yamamoto, Methanol oxidation catalysis and substructure of PtRu bimetallic nanoparticles, Applied Catalysis A: General, 326(2) (2007) 194-201. [91] B. Fang, M. Kim, J.-S. Yu, Hollow core/mesoporous shell carbon as a highly efficient catalyst support in direct formic acid fuel cell, Applied Catalysis B: Environmental, Available online (2008). [92] F. Charreteur, F. Jaouen, S. Ruggeri and J.-P. Dodelet, Fe/N/C non-precious catalysts for PEM fuel cells: Influence of the structural parameters of pristine commercial carbon blacks on their activity for oxygen reduction, Electrochimica Acta, 53(6) (2008) 2925-2938. [93] A. Garsuch, R. d'Eon, T. Dahn, O. Klepel, R.. R.. Garsuch, and J. R. Dahn, JES, 2008, 155(3) B236. [94] T. Pichler, H. Rauf, M. Knupfer, J. Fink and H. Kataura, A photoemission study of the nature of the metallic state in single wall carbon nanotube bundles at low potassium doping, Synthetic Metals, 153 (1-3) (2005) 333. [95] H. Shioyama, Y. Yamada, A. Ueda and T. Kobayashi, Graphite intercalation compounds as PEMFC electrocatalyst supports, Carbon, 43(11) (2005) 2374. [96] R. Bashyam and P. Zelenay, A class of non-precious metal composite catalysts for fuel cells, Nature, 443(7) (2006) 63. [97] G. Faubert, R. Cote, J.P. Dodelet, M. Lefevre, P. Bertrand, Oxygen reduction catalysts for polymer electrolyte fuel cells from the pyrolysis of FeII acetate adsorbed on 3,4,9,10-perylenetetracarboxylic dianhydride, Electrochim, Acta 44 (1999) 2589. [98] M. Lef`evre, J.P. Dodelet, P. Bertrand, Molecular Oxygen Reduction in PEM Fuel Cells: Evidence for the Simultaneous Presence of Two Active Sites in Fe-Based Catalysts, J. Phys. Chem. B, 106 (2002) 8705. [99] C. Bianchini, G. Giambastiani, I. G. Rios, G. Mantovani, A. Meli and A.Segarra, Ethylene oligomerization, homopolymerization and copolymerization by iron and cobalt catalysts with 2,6-(bis-organylimino)pyridyl ligands, Chemistry Reviews, 250 (11-12) (2006) 1391-1418. [100] Cicero W.B. Bezerra, Lei, A review of Fe–N/C and Co–N/C catalysts for the oxygen reduction reaction, Electrochimica Acta, 53 (15) (2008) 4937-4951. [101] R. Jasinski, A New Fuel Cell Cathode Catalyst, Nature, 201 (1964) 1212-1213. [102] A.A. Lavrentyev, B.V. Gabrelian, V.B. Vorzhev, I.Ya. Nikiforov, O.Yu. Khyzhun, Electronic properties of cubic TaCxN1 − x: A comparative study using self-consistent cluster and ab initio band-structure calculations and X-ray spectroscopy Journal of Alloys and Compounds, 472(1-2) (2009) 104-111. [103] A.L. Bouwkamp-Wijnoltz, W. Visscher, J.A.R. Van Veen and S.C. Tang, Electrochemical reduction of oxygen: an alternative method to prepare active CoN4 catalysts, Electrochimica Acta, 45 (3) (1999) 379-386.

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[104] A. Ishihara, S. Doi, S. Mitsushima, K. Ota, Tantalum (oxy)nitrides prepared using reactive sputtering for new nonplatinum cathodes of polymer electrolyte fuel cell, Electrochimica Acta, 53 (16) (2008) 5442-5450. [105] Y. Hakuta, H. Hayashi, K. Arai, Fine particle formation using supercritical fluids, Current Opinion in Solid State and Materials Science, 7( 4-5) (2003) 341-351. [106] W. Mormann, H. Jung, D. Spitzer , Ammonia as reagent or reaction medium for polymers, Super-critical Fluids as Solvents and Reaction Media, (2004) 593-616. [107] O. Merdrignac-Conanec, K. E. Badraoui, P. L’Haridon , Nitridation under ammonia of high surface area vanadium aerogels, Journal of Solid State Chemistry, 178(1) (2005) 218-223. [108] A. Leineweber, H. Jacobs, S. Hull, Ordering of Nitrogen in Nickel Nitride Ni3N

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[109] S. Desmoulins-Krawiec, C. Aymonier, A. Loppinet-Serani, F. Weill, S. Gorsse, J. Etourneau, F. Cansell, Synthesis of nanostructured materials in supercritical ammonia: nitrides, metals and oxides, J. Mater. Chem, 14 (2004) 228-232. [110] S. Slade, S. A. Campbell, T. R. Ralph, and F. C. Walsh Ionic Conductivity of an Extruded Nafion 1100 EW Series of Membranes, J. Electrochem. Soc. 149 (2002) A1556-A1564. [111] L. Merlo, A. Ghielmi, L. Cirillo, M. Gebert and V. Arcella, Resistance to peroxide degradation of Hyflon® Ion membranes, Journal of Power Sources, 171(1) (2007) 140-147. [112] I.H. Hristov, S.J.Paddison, R. Paul, Molecular dynamics simulations of proton diffusion in the short-side-chain perfluorosulfonic acid ionomer, ECS Transactions 11 (1) (2007) 789-795. [113] M.T. Hicks, DOE Hydrogen program report: MEA and Stack Durability for PEM Fuel Cells, Contract No. DE-FC36-03GO13098. [114] R. Souzy, B. Ameduri, B. Boutevin, G. Gebel, P. Capron, Functional fluoropolymers for fuel cell membranes, Solid State Ionics, 176 (39-40) (2005) 2839-2848. [115] J. D. Weaver, E. L. Tasset and W.E. Fry, Supported fluorocarbonsulfonic acid catalysts, Catal. Today, 14 (1992) 195–210. [116] J. Chlistunoff, F. Uribe, B. Pivovar, B., Oxygen reduction at the Pt/recast-Nafion® film interface. Effect of the polymer equivalent weight, ECS Transactions, 2 (8) (2007) 3746. [117] I. H. Hristov, S. J. Paddison, R. Paul, Molecular modeling of proton transport in the short-side-chain perfluorosulfonic acid ionomer, Journal of Physical Chemistry B 112 (10), (2008 ) 2937-2949. [118] G. Li and P. G. Pickup, Ionic Conductivity of PEMFC Electrodes, J. Electrochem. Soc. 150 (2003) C745-C752. [119] H..Xu, H. R. Kunz, J. M. Fenton, and L. J. Bonville, Improvement of PEM Fuel Cell Performance Using Low Equivalent Weight PFSA Ionomers, ECS Trans. 3(1) 361 (2006) [120] D.E. Curtin, R.D. Lousenberg, T.J. Henry, P.C. Tangeman, M.E. Tisak, Advanced materials for improved PEMFC performance and life, Journal J. Power Sources, 131(1/2) (2004) 41-48.

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[121] I. Gatto, A. Saccà, A. Carbone, R. Pedicini, F. Urbani, E. Passalacqua, CO-tolerant electrodes developed with PhosphoMolybdic Acid for Polymer Electrolyte Fuel Cell (PEFCs) application, Journal of Power Sources, l171( 2) (2007) 540-545. [122] H. Xu, Y. Song, H.R. Kunz, and J.M. Fenton, J. of the Electrochem. Soc., 152 (2005) A1828-A1836. [123] I. V. Kozhevnikov, Heterogeneous acid catalysis by heteropoly acids: Approaches to catalyst deactivation, J. Molecular Catalysis A: Chemical, (2009) Availavle online. [124] V. Ramani, S. Swier, M.T. Shaw, R.A. Weiss, H.R. Kunz, J.M., Fenton, Membranes and MEAs Based on Sulfonated Poly(ether ketone ketone) and Heteropolyacids for Polymer Electrolyte Fuel Cells Journal of the Electrochemical Society, 155 (6) (2008) B532-B537. [125] V. Ramani, A. Smirnova, H.R. Kunz, J.M. Fenton, Performance of high temperature proton exchange membrane fuel cells under off-design conditions (2005) Proceedings of Electrochemical Society, 31 (2002) 426-439. [126] H. Xu, H.R. Kunz, L.J. Bonville, and J.M. Fenton, J. of the Electrochem. Soc., 154 (2007) B271-B278. [127] V. Ramani, H.R. Kunz, J.M. Fenton, Effect of particle size reduction on the conductivity of Nafion®/phosphotungstic acid composite membranes Journal of Membrane Science 266 (1-2) (2005) 110-114. [128] http://en.wikipedia.org/wiki/Keggin_structure [129] U. Lavrenic Stangar, N. Groselj, B. Orel, and Ph. Colomban, Structure of and Interactions between P/SiWA Keggin Nanocrystals Dispersed in an Organically Modified Electrolyte Membrane, Chem. Mater., 12 (12), 3745 -3753, 2000. H. Xu, H.R. Kunz, L.J. Bonville, and J.M. Fenton, Jl. of the Electrochem. Soc., 154, B271-B278 (2007). [130] Z.-G.Shao, H.Xu, M.. Li, I.-M.Hsing, Hybrid Nafion-inorganic oxides membrane doped with heteropolyacids for high temperature operation of proton exchange membrane fuel cell, Solid State Ionics 177 (7-8) (2006) 779-785. [131] R. J. Stanis, M.-C. Kuo, A. J. Rickett, J. A. Turner, Andrew M. Herring, Investigation into the activity of heteropolyacids towards the oxygen reduction reaction on PEMFC cathodes, Electrochimica Acta, Available online (2008).

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Chapter 6

AIR-BREATHING DIRECT METHANOL FUEL CELLS WITH CATALYSED TITANIUM MESH ELECTRODES Raghuram Chetty1 2* and Keith Scott1 1

School of Chemical Engineering & Advanced Materials, Newcastle University, Newcastle upon Tyne, NE1 7RU, United Kingdom. 2 Department of Chemical Engineering, Indian Institute of Technology Madras, Chennai, 600036, India.

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ABSTRACT The conventional electrode structure used in the DMFC is generally based on porous carbon-based gas diffusion electrodes, which consists of a catalyst layer, hydrophobic microporous layer and carbon paper or cloth. This structure is not ideal for transport and release of CO2 gas produced during methanol oxidation at the anode, potentially resulting in considerable hydrodynamic and mass transport limitations. To circumvent these problems, an expanded titanium mesh has been adopted as the catalyst substrate material in this study. Titanium mesh was used as the substrate due to its chemical stability and its ability to support a diverse range of electrocatalysts. In the proposed fuel cell application, the mesh-based electrode has several potential advantages in terms of cost, simplicity, size and shape. This work describes the design, fabrication and evaluation of a passive air-breathing direct methanol fuel cell using the mesh-based electrodes. PtRu/Ti and Pt/Ti prepared by electrodeposition onto the Ti mesh were used as anode and cathode, respectively. Methanol is stored in an in-built reservoir and oxygen is taken from the surrounding air. Single cells with an active area of 9 cm2 produced a power density of 9.5 mW cm-2, whereas a cubic four-cell stack with an active area of 24 cm2 produced a maximum power of 180 mW. The effects of experimental parameters such as concentration and temperature on the cell performance were also investigated.

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1. INTRODUCTION Direct methanol fuel cells (DMFC) are considered as a key energy source for portable applications given the relatively simple system design, cell operation as well as easy and instantaneous refueling [[1] ,[2] ]. DMFC that uses liquid methanol directly without a reformer has a higher specific energy density (3000 Wh kg-1) than the lithium-ion battery system (200 Wh kg-1) [[3] ] and can be operated at ambient conditions. These advantages make DMFC a potential candidate for portable electronic applications such as mobile phones, notebooks, personal digital assistance (PDAs), etc whose power requirement have fueled the need for high energy density power sources. However, further improvements in the performance of the DMFC are required before its practical applications can be fully realized, and thus numerous amount of research has gone into their development [[4] -[7] ]. Increasing the catalyst efficiency, improving the fuel cell design and optimizing the structure of the electrode are some of the key strategies adopted in the recent fuel cells research and development. Air-breathing direct liquid fuel cells, which operates without the need for external pumps or other ancillary devices for fuel and oxidant has attracted great interest since it reduces the volume, costs and energy consumption of the fuel cell system [[8] -[18] ]. Nevertheless, passive direct liquid fuel cells have more challenges with respect to mass transport of reactant and products [[16] ]. The DMFC generally suffers from mass transport limitations at the anode, due to the low diffusion coefficient of methanol in water and the release of carbon dioxide gas bubbles during methanol oxidation (CH3OH + H2O → CO2 + 6H+ + 6e−). A methanol concentration gradient also exists within the thickness of the catalyst layer, which results in poor utilization of the catalyst [[19] ]. The conventional electrode structure used in the DMFC are generally based on porous carbon-based gas diffusion electrodes, which consists of a catalyst layer, microporous layer and carbon paper or cloth impregnated with carbon black and hydrophobic polymer. This structure is not ideal for transport and release of carbon dioxide gas from the anode, potentially resulting in considerable hydrodynamic and mass transport limitations [[20] ]. Recently we have reported the use of a mesh-based electrode; a hydrophilic anode coated with platinum based catalyst for direct liquid fuel cells [[21] -[24] ], which showed promising behavior in terms of gas removal and electrical performance in comparison to the conventional carbon supported membrane electrode assembly (MEA). Titanium mesh was used as the substrate due to its chemical stability and its ability to support a diverse range of electrocatalysts. The use of fine metal mesh eliminates the need for gas diffusion electrode and increases access of methanol to the catalyst layer. Moreover, the mesh substrate has advantages of simplicity in catalyst coating, low cost fabrication and flexibility in terms of shape and size of fuel cells. Electrodes of metal and metal oxide coatings on titanium are widely used in a variety of industrial electrochemical processes and procedures for catalyst deposition are well known; including thermal, electrochemical and vapor phase deposition [[25] ] In this work, we report on some of our recent activities regarding the research and development of air-breathing DMFC operating at room temperature using electrodeposited Pt-based catalyst on titanium-mesh electrodes, which are intended for portable electronics applications.

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2. EXPERIMENTAL The electrodes were prepared by electrodeposition of Pt or PtRu catalyst onto titanium mesh. The titanium substrate (2Ti5-031 obtained from Dexmet Corporation, USA) was boiled in concentrated hydrochloric acid for 2 minutes and washed with water. The treated surface was then immersed in a boiling aqueous solution of oxalic acid (15 %) and finally washed with water and stored in water until use. For electrodeposition, the pretreated Ti mesh was immersed in an aqueous 0.01 M H2SO4 solution containing 2.0 mM H2PtCl6 + 2.0 mM RuCl3 (Alfa Aesar). Electrodeposition was performed at room temperature at a constant current density of 4 mA cm-2 [[20] ]. The desired catalyst loading was achieved by controlling the electrodeposition time. The electrodes prepared were sonicated with water and the amount of catalyst deposited was determined by comparing the substrate weight before and after deposition. The electrodes were then characterised by Scanning Electron Microscopy (SEM, Philips XL30 ESEM-FEG) interfaced with energy dispersive X-ray analysis (EDX).

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Membrane Electrode Assembly The electrodeposited PtRu/Ti mesh was used as anode, with metal loading of 4 mg cm-2 PtRu. The cathode was electrodeposited Pt/Ti mesh with a loading of 4 mg cm-2 Pt. The membrane electrolyte used in this work was pretreated Nafion® 117. The pre-treatment involved boiling the membrane for 1 h in 5 vol. % H2O2 and 1 h in 0.5 M H2SO4 before washing in boiling deionised water. The electrodes were coated with 1 mg cm-2 of Nafion® solution prior to forming a MEA. Finally the MEA was obtained by pressing the catalystcoated cathode and anode on either side of the pre-treated Nafion® -117 proton exchange membrane by hot pressing at 130 °C and 100 kg cm-2 for 3 minutes [[21] ]. Figures 1 show the structure of the mesh based MEA in comparison to the conventional carbon based MEA. One of the main challenges with direct liquid fuel cells is that the anode has to facilitate two-phase mass transfer, i.e. liquid fuel transport and gaseous product evolution transport. In the conventional carbon based electrodes, generally hydrophilic micropores in the diffusion electrode are required to transport liquid fuels, and hydrophobic micropores are favourable for gas transport. Polytetrafluoroethylene (PTFE) and Nafion® are the two commonly used polymeric additives used mainly as the bonding agent in the polymer electrolyte membrane fuel cell electrodes. The presence of PTFE tends to make the electrode structure more hydrophobic, whereas Nafion® makes the electrode more hydrophilic. The typical conventional electrode consists of three parts: backing layer, gas diffusion layer and catalyst layer. Teflonised carbon paper or cloth is generally used as the backing layer, which is approximately 250-350 μm thick. The gas diffusion layer consists of un-catalysed carbon, with Nafion® or Teflon with 30-50 μm thick, and the catalyst layer consisting of supported or unsupported Pt-based catalyst and Nafion® of 20 to 50 μm thick. In contrast, the mesh-based MEA was thinner; composed of a 50 μm thick Ti mesh and a 3-5 μm catalyst layer. PTFE was not employed in the anode structure and hence it is more hydrophilic than the conventional electrodes.

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(a)

(a)

(e)

(b) (c) (d)

(i)

(ii)

Figure 1. Schematic representation of (i) conventional and (ii) mesh-based anode. (a) Polymer electrolyte membrane, (b) catalyst layer, (c) microporous layer, (d) backing layer and (e) catalyst coated titanium mesh.

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Single Cell Fixture The prepared MEA was fitted into a small fixture made of transparent acrylic, as shown in Fig. 2, to measure the single cell performance under air-breathing conditions. The catalyst coated mesh act as the current collector for the anode and cathode side respectively. Silicone rubber was used as the gasket at the perimeter of the MEA and acrylic fixture, and the cell was held together with several threaded tie-bolts with nuts. The active area of the electrode was 9 cm2, and with this configuration, the cell operates without any external devices for feeding methanol or blowing air into the system. Oxygen can diffuse into the cathode from the ambient due to the air-breathing action of the cell [[10] ]. Similarly, methanol diffuse to the electrode available from the reservoir built in the anode fixture. The performance of the cell was measured after it was kept in deionised water for overnight prior to the test. Polarisation data were measured manually using a Kenwood PEL-151 electronic load, a 30 second waiting time was used to obtain stable voltage.

Design of Stack The MEA for four-cell stack were made and then fitted into a cubic acrylic fixture. A picture of the cubic air-breathing DMFC cell pack is shown in Fig. 3, each MEA has an active area of 6 cm2, giving a total of 24 cm2 for the cubic cell. The outer dimensions of the stack are 40mm x 40mm x 50mm. The cathodes on all the sides of the cubic pack are exposed to

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ambient air and methanol solution is stored between the MEAs. The stack was assembled by connecting the anode and cathode in series. Other features were similar with the single cell as explained above. Si rubber Membrane

Methanol reservoir

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Pt/Ti cathode

50 mm PtRu/Ti anode

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Figure 2. (a), Schematic representation of the single air-breathing mesh-based DMFC and (b) photograph of the fabricated cell.

Methanol reservoir

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Membrane

50 mm

40 mm MEA

(a)

Air

(b)

Figure 3. The cubic air-breathing mesh-based direct methanol fuel cell stack. (a) Design and (b) fabricated. Anode: electrodeposited Pt/Ti, cathode: electrodeposited PtRu/Ti.

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3 RESULTS AND DISCUSSION To examine the morphology and composition of the catalyst coated mesh, SEM and EDX were employed. A representative SEM image showing electrodeposition PtRu/Ti electrode is shown in Fig. 4(i). As can be seen from the figure, the electrodeposited catalyst layer is quite rough in morphology due to the porosity induced by gas evolution during electroplating. Fig. 4(ii) shows the SEM at higher magnification and EDX mapping. The EDX analysis showed that the atomic ratio of Pt:Ru is 1:1. As can be seen from the elemental distribution images, Pt and Ru are uniformly distributed over the titanium mesh confirming the homogeneous catalysts coating by electrodeposition.

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(i)

(a)

(b)

(c)

(d)

(ii)

Figure 4. (i) Scanning electron micrograph (bar = 100 μm) of electrodeposited PtRu on Ti mesh and (ii) SEM/EDX mapping showing the distribution of Pt and Ru on the Ti mesh electrode: (a) SEM image at higher magnification (bar = 5 μm), EDX mapping of (b) Pt, (c) Ru and (d) Ti.

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Figure 5. Fuel cell polarisation curves of air-breathing DMFC single cell using catalystcoated titanium mesh electrodes with various concentration of methanol. Anode catalyst: 4 mg cm-2 PtRu and cathode catalyst: 4 mg cm-2 Pt. Current density (open symbol) and power density (blocked symbol) for 1 mol dm-3 (V,T), 2 mol dm-3 (Δ,c), 4 mol dm-3 (○, ●) and 5 mol dm-3 (ϒ, ■) methanol. The effect of experimental parameters on the single cell performances under air-breathing conditions was investigated. The concentration of methanol at the anode has a major influence on the operation of the air-breathing DMFC. As reported in the literature, high concentrations are known to cause an increase in methanol crossover. Optimisation is required between methanol oxidation kinetics and methanol crossover in order to enhance the performance of the DMFC. Fig. 5 shows the effect of methanol concentrations, from 1 to 5 mol dm-3, on the single cell performance. At higher methanol concentrations, the open circuit voltage (OCV) decreased: OCV’s with 1, 2, 4 and 5 mol dm-3 methanol were 0.647, 0.615, 0.602 and 0.578 V, respectively. The decrease in OCV reflects the characteristic of increased methanol crossover at higher concentrations [[26] ]. As can be seen from the figure (Fig. 5), power densities of the cell increased with increasing methanol concentrations up to 4 mol dm3 and decreased with further increase in concentration, exhibiting a maximum power density of 9.5 mW cm-2. The optimum methanol concentration was found to be higher than that reported for conventional liquid and air feed DMFCs, which has a maximum performance at methanol concentrations of 1-2 mol dm-3 [[27] , [28] ]. However, under air-breathing conditions, 1 mol dm-3 methanol showed a poor performance and the voltage fell at higher current densities (i.e. > 30 mA cm-2). On the other hand, concentrations higher than 4 mol dm3 methanol could facilitate methanol crossover through the membrane and reduce cell performance through a mixed potential polarising the cathode [[29] ]. In a DMFC, CO2 is produced at the anode by oxidation of methanol. In this study, transparent acrylic plates were used as cell fixtures and thus CO2 bubbles could be observed at the anode and rising by the buoyancy force [[30] ] (see Fig. 6). The use of mesh based electrodes assists the release of CO2 bubbles from the anode and reduces the negative effect of the bubbles restricting diffusion of methanol to the active catalytic sites. Under airbreathing conditions, water also accumulates at the cathode side to limit access of oxygen to

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the catalytic sites. Therefore, the use of mesh based electrodes has the potential advantage of facilitating removal of CO2 bubbles and water from the air-breathing fuel cells.

Figure 6. CO2 bubbles generated at the anode side of the air-breathing single cell.

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Cell voltage / V

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An attempt was made to evaluate the performance of a cubic four-cell stack, in which the four MEAs were electrically connected in series to meet the voltage requirement of power conditioning for portable electronic gadgets. The cathodes on all the sides of the cubic stack were exposed to ambient air and 40 ml methanol solution was stored in the reservoir between the MEAs. The power performance of the four-cell stack with 2 and 4 mol dm-3 methanol concentrations is shown in Fig. 7. The OCVs were 2.188 and 2.156 for 2 mol dm-3 and 4 mol dm-3 methanol, respectively and the limiting current densities were almost same for both the concentrations. Similar to the results with single cells, the stack performance also increased with increasing methanol concentration; in the present case the cell power increased from 148 mW with 2 mol dm-3 to 180 mW with 4 mol dm-3 methanol concentration. The decrease in power performance observed for lower methanol concentrations could be due to the mass transport limitation, as explained before.

0 400

Current / mA

Figure 7. Performance of the cubic air-breathing DMFC four-cell stack with an active area of 24 cm2 fabricated using catalyst-coated titanium mesh electrodes. Fuel cell polarisation plots for 2 mol dm-3 (□, ■) and 4 mol dm-3 (○, ●) methanol concentration. Anode catalyst: 4 mg cm-2 PtRu and cathode catalyst: 4 mg cm-2 Pt.

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o

Temperature / C

40

35

30

25

20

0

100

200

300

Current / mA Figure 8. Variation in cell operating temperature of the cubic air-breathing four-cell DMFC stack.

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In stack operation, the temperature of fuel cell generally increased due to the heat released during electrochemical reaction. During fuel cell polarisation, the temperature of the cell was monitored by a digital thermometer connected to a small thermocouple inserted at the anode. Fig 8 shows the temperature change in the air-breathing DMFC stack with respect to current produced for methanol concentrations of 2 and 4 mol dm-3. The cell temperature at 4 mol dm-3 methanol concentration was higher than that at 2 mol dm-3 methanol concentration: in the case of the cell using 2 mol dm-3 methanol, the cell temperature rose from 23 °C up to 33 °C, whereas with the 4 mol dm-3 methanol the temperature rose to around 36 °C. This suggests that higher methanol concentration could also enhance the performance of the fuel cells due to the increase in operating temperature.

4. SUMMARY AND CONCLUSION This study presents the results of air-breathing DMFC, fabricated using electrodeposited Pt cathode and PtRu anode on titanium mesh-based substrates. The DMFC single cell with an active area of 9 cm2 under air-breathing conditions produced a maximum performance of 9.5 mW cm-2 with a 4 mol dm-3 methanol concentration. Since the transport of aqueous methanol fuel and the oxidation product CO2 gas can rapidly pass through the pores of mesh-based electrodes, the electrochemical reaction at the catalyst sites can proceed rapidly with less mass transfer limitations and increased catalyst utilization. A cubic four-cell stack with a total active area of 24 cm2 behaved in a similar way to the single cell with respect to the effect of methanol concentration and produced a power of 180 mW. The attractions of using meshbased electrode are the simplicity of electrocatalyst preparation by electrodeposition and flexibility of the mesh, which can result in variety of shapes. Further research and development on the mesh-based electrodes could significantly improve power and energy density, efficiency and reliability of the air-breathing fuel cells for portable applications.

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ACKNOWLEDGMENTS The authors thank the United Kingdom Engineering and Physical Sciences Research Council (EPSRC) for financial support, and Dexmet Corporation, USA for providing titanium mesh.

REFERENCES [1] [2] [3] [4] [5] [6] [7] [8] [9] [10]

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[11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25] [26]

Dillon, R.; Srinivasan, S.; Aricò, A. S.; Antonucci, V. J. Power Sources 2004, 127, 112-126. Scott, K.; Shukla, A. K. Modern Aspects of Electrochemistry; White, R.E. (Ed.) Springer: 2007; Vol. 40, pp. 127-227. Han, J.; Park, E. S. J. Power Sources 2002, 112, 477-483. Neburchilov, V.; Martin, J.; Wang H.; Zhang, J. J. Power Sources 2007, 169, 221-238. Liu, H.; Song, C.; Zhang, L.; Zhang, J.; Wang, H.; Wilkinson, D. P. J. Power Sources 2006, 155, 95-110. DeLuca, W. N.; Elabd, Y. A. J. Polym. Sci. Part B: Polym. Phys. 2006, 44, 2201-2225. Antolini, E. Mat. Chem. Phys. 2003, 78, 563-573. Faghri, A.; Guo, Z. Appl. Therm. Eng. 2008, 28, 1614-1622. Yang, T.; Shi, P. Int. J. Hyd. Energy 2008, 33, 2795-2801. Baglio, V.; Stassi, A.; Matera, F. V.; Di Blasi, A.; Antonucci, V.; Arico, A. S. J. Power Sources 2008, 180, 797-802. Chen, R.; Zhao, T. S. Electrochem. Commun. 2007, 9, 718-724. Chan, Y. H.; Zhao, T. S.; Chen, R.; Xu, C. J. Power Sources 2008, 178, 118-124. Song, K.Y.; Lee, H. K.; Kim, H. T. Electrochim. Acta 2007, 53, 637-643. Martin, J. J.; Qian, W.; Wang, H.; Neburchilov, V.; Zhang, J.; Wilkinson, D. P.; Chang, Z. J. Power Sources 2007, 164, 287-292. Pan, Y.H. J. Power Sources 2006, 161, 282-289. Qian, W.; Wilkinson, D. P; Shen, J.; Wang, H.; Zhang, J. J. Power Sources 2006, 154, 202-213. Liu, J.; Zhao, T. S.; Chen, R.; Wong, C. W. Fuel Cells Bulletin 2005, 12-17. Chen, C. Y.; Yang, P. J. Power Sources 2003, 123, 37-42. Shao, Z. G.; Zhu, F.; Lin, W. F.; Christensen, P. A.; Zhang, H. Phys. Chem. Chem. Phys. 2006, 8, 2720-2726. Allen, R. G.; Lim, C.; Yang, L. X.; Scott K.; Roy, S. J. Power Sources 2005, 143, 142149. Chetty, R.; Scott, K. Electrochim. Acta 2007, 52, 4073-4081. Chetty, R.; Scott, K. J. Power Sources 2007, 173, 166-171. Chetty, R.; Scott, K. J. New Mat. Electrochem. Sys. 2007, 10, 135-142. Chetty, R.; Scott, K. J. Appl. Electrochem. 2007, 37, 1077-1084. Trasatti, S. Electrochim. Acta 2000, 45, 2377-2385. Liu, J. G.; Zhao, T. S.; Chen, R.; Wong, C. W. Electrochem. Commun. 2005, 7, 288294.

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[27] Gurau, B.; Smotkin, E. S. J. Power Sources 2002, 112, 339-352. [28] Valbuena, W. H. L.; Paganin, V. A.; Gonzalez, E. R. Electrochim. Acta 2002, 47, 37153722. [29] Bae, B.; Kho, B. K.; Lim, T. H.; Oh, I. H.; Hong, S. A.; Ha, H. Y. J. Power Sources 2006, 158, 1256-1261. [30] Argyropoulos, P.; Scott, K.; Taama, W. M. Electrochim. Acta 1999, 44, 3575-3584.

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In: Polymer Electrolyte Membrane Fuel Cells… Editors: R. Esposito

ISBN: 978-1-60692-773-1 ©2009 Nova Science Publishers, Inc.

Chapter 7

METAL ELECTROCATALYSTS FOR DIRECT LIQUID-FEED FUEL CELLS Umit B. Demirci* Université Lyon 1, CNRS, UMR 5615, Laboratoire des Multimatériaux et Interfaces, 43 boulevard du 11 Novembre 1918, F-69622 Villeurbanne, France.

ABSTRACT

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Direct liquid-feed fuel cells (DLFCs) that are intended for mobile/portable devices are technologies based on the polymer electrolyte membrane fuel cell (PEMFC). The technical, commercial development of DLFCs is dependent on various materials like, e.g., metal electrocatalysts. These electrocatalysts are key materials. They catalyze the fuel oxidation as well as the oxygen reduction. The main difference between DLFCs and PEMFC is the fuel, that is, liquid chemicals and hydrogen gas respectively. The early investigations about DLFCs assessed Pt electrocatalysts that were being used for the PEMFC but with the fuels being different, new issues arose. For instance, in the direct methanol fuel cell (DMFC), Pt poisoned because of the formation of CO-like intermediates like CO, formaldehyde or formic acid. Hence new metal-based catalytic materials were regarded and are today being investigated. This is discussed in the present chapter. Many liquid chemicals have been envisaged as fuels of DLFCs but the paper only deals with the main ones: i.e., the DLFCs fed with, e.g., C1-C3 alcohols, formic acid and sodium borohydride. In the DLFCs, whatever the fuel (e.g., methanol, ethanol, formic acid, sodium borohydride...) is, the anode electrocatalysts are generally metal-based. Monometallic electrocatalysts have been widely investigated and, today, lots of new bimetallic (and even trimetallic) combinations are suggested throughout the open literature. Globally, the investigations devoted to the DLFCs electrocatalysts focus on reactivity and durability. With the multimetallic electrocatalysts, improved reactivity and durability may be obtained because of electronic and/or geometric effects. In other words, the addition of a metal (e.g., Ru) to another (e.g., Pt) may modify the electronic and geometric features of the latter, influencing then its adsorption properties, reactivity and durability. This, especially, is discussed in the present chapter. *

Tel. +33 472 44 84 03 Fax. +33 472 44 06 18 Mail. [email protected]

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Umit B. Demirci Theoretical results/data (e.g., d-band centre of metals and segregation, both from density functional theory) are useful tools for trying to understand the catalytic behaviors of metal electrocatalysts. These tools, available in the open literature, are used in the present discussion for interpreting electronic and geometric effects. Furthermore, when possible these are exploited for suggesting alternative multimetallic combinations as potential anode electrocatalysts for the DLFCs that are tackled.

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1. INTRODUCTION Hydrogen and fuel cells are technologies whose recent scientific, technical developments highly benefit from difficult contexts in energy and the environment. Today the world is completely energy-dependent and energy has become a major worldwide stake. Energy is very largely fossil fuels-based and the resources are depleting. In such a critical energy context, there is unanimity for finding solutions. Parallel to that, environmental problems have become a major worldwide concern. Although many environmental problems face our planet, global warming (climate change) seems to be the most important, likely because it is manifestly the most disastrous as well as appearing in most of the media. It is commonly recognized that anthropogenic greenhouse gases (emitted by human activities) are mainly responsible for global warming. The greenhouse effect is a natural phenomenon regulating the earth’s temperature but the anthropogenic greenhouse gases cause an additional greenhouse effect implying global warming and the consequent climate change. Several kinds of greenhouse gases exist, among which two gases stand out. The first one, and also the main one, is carbon dioxide, which mainly stems from the combustion of fossil fuels. The second one, methane, has many sources; one of them is fossil fuel production (i.e., coal mining, natural gas and petroleum systems). Globally, all of these energy- and pollution-related problems favor the development of hydrogen. Molecular hydrogen is today regarded as one of the most promising alternatives to fossil fuels. It is a highly energetic, non toxic compound but it is highly flammable and explosive. Hydrogen is abundant on earth but exists combined with other atoms like, e.g., water (H2O), methane (CH4) or cellulose ((C6H10O5)n). Hence it has to be produced. This is one of the main issues holding up the development of a hydrogen economy [1]. The second main issue is related to the gaseous nature of molecular hydrogen, which is the lightest compound, which implies significant storage problems [2]. Investigations are today in progress in order to solve these production and storage problems. With respect to the production, about ten processes are being studied. These processes are based on three main methods, which are thermal, electrochemical and biological. These processes convert three categories of raw materials: fossil fuels, biomass and water. Regarding the storage, six main hydrogen storage methods are in development: high pressure tanks, liquid hydrogen, adsorption of hydrogen on materials with a large specific surface area, hydrogen adsorption on interstitial sites in a host metal, chemically bonded in covalent and ionic compounds, and oxidation of reactive metals (e.g., Li, Na, Mg, Al, Zn) with water [2]. Hydrogen is an energy carrier and in that sense it is intended to fuel an energy converter, which is none other than the fuel cell. Basically, hydrogen is today the main fuel of the fuel cells. Note that hydrogen can also be regarded as an energy source when it directly feeds an engine.

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Fuel cells are power generators whose principle is quite simple: they convert chemical energy (e.g., that contained in molecular hydrogen) directly into electrical energy (Figure 1). They are particularly attractive because of their efficiency and the environmental benefits associated with their utilization. About this latter advantage, the hydrogen-powered fuel cell is generally regarded to be a green technology, that is, in accordance with green chemistry. Green chemistry is an emerging concept looking to promote innovative chemical technologies that reduce or eliminate the use or generation of hazardous substances in the design, manufacture and use of chemical products [3, 4]. Green chemistry is engaged in addressing energy needs through the development of more sustainable energy technologies, like, e.g., fuel cells. Renewable resource utilization, like, e.g., hydrogen from water or biomass, is a central tenet of green chemistry [3]. Using hydrogen-powered fuel cells compared to more conventional sources of energy would provide four major potential benefits: an overall efficiency higher than that of the conventional heat engine, pollution minimized because of the generation of only water as a by-product, the possibility of obtaining hydrogen (i.e., the fuel) from renewable resources, and its higher reliability [4].

2 eInlet: Fuel H2

Outlet: H 2O

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2 H+

On Fig. 1, the words « Anode », « Cathode », « Polymer Electrolyte Membrane » have been cut.

Anode

2 H+

Polymer Electrolyt eMembra

Inlet: ½ O2 Cathod

Figure 1. Hydrogen-powered PEM-type fuel cells.

Several types of fuel cells mainly differing from the nature of their electrolyte are under active development: for instance, the polymer electrolyte membrane fuel cell (PEMFC), the alkaline fuel cell (AFC) or the solid oxide fuel cell (SOFC) [5]. The hydrogen-powered PEMFC is a low-temperature fuel cell that is intended for portable, mobile and automotive applications. However, issues related to hydrogen storage have favored the emergence of alternative fuels, specifically liquid fuels, for the PEMFC intended for portable and mobile applications. Methanol is the most investigated liquid fuel. The methanol-powered PEMFC is the direct methanol fuel cell (DMFC) whose commercial potential is commonly considered as being significant and promising [5, 6]. Compared to hydrogen, methanol is easily handled, transported and stored, even if it is toxic, flammable and not a primary fuel. Besides methanol, lots of chemicals have in fact been assessed as a potential liquid fuel for the PEMFC. They are ammonia, borohydride aqueous alkaline solution, dimethoxymethane,

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dimethyl ether, ethanol, ethylene glycol, formic acid, hydrazine, 1-methoxy-2-propanol, 1propanol, 2-propanol, tetramethyl orthocarbonate, trimethoxymethane, trioxane [7]. These liquid chemical-powered fuel cells are called direct liquid-feed fuel cells (DLFCs) (Figure 2).

y e(a)

On Fig 2a) and b), the words , , , have been cut On Fig 2 b), The first arrow should contain the text

Inlet: Liquid Fuel

Outlet: y/2 H2O y H+

Outlet: x CO2 Anode

Polymer Electrolyt eMembra

Inlet: y/4 O2 Cathod

6 e(b) Inlet: CH3OH + HO

Outlet: 3 H 2O 6 H+

Outlet: CO2

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Anode

Polymer Electrolyt eMembra

Inlet: 3/2 O2 Cathod

Figure 2. Direct liquid-feed fuel cells (a) and direct methanol fuel cell (b).

As was remarked by Steele and Heinzel [8], before PEMFC as well as DLFCs can gain a significant share of the electrical power market, important issues have to be addressed, one of them is associated with the anode electrocatalyst. Steele and Heinzel [8] has reported that for hydrogen-powered PEMFC, Pt is the most active metal but because it is one of the most expensive metals the reduction in Pt content in order to reduce the cost without degrading the cell performance is an important research activity. The utilization of Pt requires high purity hydrogen, which means hydrogen without any trace of CO, which is known to be an effective poison of pure Pt. With respect to the DMFC, the situation is quite similar. For the methanol electrode, a CO-tolerant electrocatalyst is required, as adsorbed CO-like species are formed during the oxidation of the fuel. This remains one of the most challenging tasks for the successful development of a commercial fuel cell. Associating Pt with another metal is a strategy. For instance, binary Pt-based alloys were studied and the system Pt-Ru showed a CO-resistance improved against single Pt [9]. Associating Pt with another metal is also a strategy in order to reduce the cost of the electrocatalyst by reducing the Pt content [10]. Lots of bimetallic combinations, including or not Pt, are conceivable and a purely experimental

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approach should be, in this case, laborious [11]. The best approach for seeking new electrocatalysts is a combined approach of theoretical and fundamental electrochemical studies, involving knowledge of the oxidation mechanism [9]. Since CO poisoning is acute on platinum surfaces, to aid the design of more CO-tolerant anode electrocatalysts, a deeper understanding of the surface chemistry on platinum is also required [12]. In that way, density functional theory (DFT) is a theoretical, computational tool that is really efficient for predicting and explaining the improved stability of Pt-Ru in relation to pure Pt [13, 14]. Accordingly, the present chapter aims to discuss the reactivity of the metal-based anode electrocatalysts of some DLFCs, while supporting the discussion with DFT-based theoretical work that is available throughout the literature. More especially, the discussion focuses on bimetallic materials.

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2. DFT-BASED THEORETICAL CONCEPTS According to Mavrikakis, Nørskov and co-workers [15], for a class of reactions proceeding over transition metal catalysts surfaces, there is a universal, reaction-independent relationship between activation energies and the stability (binding energies) of important intermediates in the reaction. Such relationship suggests a general approach to optimizing the catalyst by searching for new materials with adsorbate-surface interaction strengths in the right range and, in that way, the DFT calculations would be an efficient, effective tool. Mavrikakis and Nørskov and co-workers are researchers deeply involved in the DFT-based calculations that are applied to mono- and bi-metallic catalytic materials. The present section reviews the main results of their work which several references are given along the chapter. Computational methods have revolutionized the approach there was for seeking new catalytic metals. The calculational methods in close conjunction with experiments can be used to develop some useful concepts to describe and understand adsorption and reactions on surfaces. The basis for the discussion is DFT [16]. DFT is a framework for the study of the electronic ground-state properties of molecules and solids [17]. According to Mavrikakis [11], a computational approach can reduce the lab work required for finding the best catalyst for a reaction. This is a first, significant advantage. Besides, in the computer experiments it is possible to study the role of electronic and geometric effects separately, what is important since both effects influences the metal catalysts reactivity [16, 18, 19]. It has especially been focused on simple concepts that can be used to classify the reactivity of transition metal surfaces, the goal being to gain an understanding of the properties of the clean surface that determine its reactivity. If these factors can be singled out, it would be possible to have powerful concepts that can be used to develop new and more effective catalysts [16].

2.1 Metal D-Band Center Hammer and Nørskov [20] demonstrated that the important surface parameters determining the reactivity are the position of the whole d-bands relative to the Fermi level, because this would determine both the size of the bonding and anti-bonding energy shifts and the degree of filling of the anti-bonding states as well as the size of the coupling matrix

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elements which also enter the size of the energy shifts. This observation is the central tenet of theoretical work that followed. Nørskov and co-workers [18] showed that an understanding of the properties of metal surfaces that determine their reactivity predicts the reactivity of a given surface, including the effect structure, the effect of alloying and the effect of coadsorbed species. In other words, the basic idea is that trends in the interaction energy between an adsorbate and a metal surface are governed by the coupling to the metal d-bands, since the coupling to the metal sp states is essentially the same for the transition and noble metals. The d-band center has been suggested. It is defined as the centroid of the d-type density of states in an atomic sphere centered at a surface atom. The d-band center is an important parameter characterizing the ability of the surface d-electrons to participate in bonding to the adsorbate. The d-band center of a given metal atom depends on the surroundings, and one of the possibilities for modifying the reactivity of a metal is by depositing it as an overlayer or alloying it into the surface layer of another metal [16, 19]. The formation of a surface metal-metal bond can produce large perturbations in the electronic, chemical, and catalytic properties of a metal. On a surface, the formation of a heteronuclear metal-metal bond induces a flow of electron density toward the element with the larger fraction of empty states in its valence band [21]. Nørskov and co-workers [18] have provided the values of the d-band centers of some transition and noble metals. Interestingly it has besides been showed how the d-band center of a metal change when a second metal is alloyed to it.

Table 1. Shifts in d-band centers of surface impurities and overlayers relative to the

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clean metal values (italic); the impurity/overlayer atoms are listed horizontally and the host entries are listed vertically; for each combination of the two numbers listed is first the isolated impurity given and than the overlayer; all values are in eV and the elemental d-band centers are relative to the Fermi level; Reprinted from [18], Copyright (1997), with permission from Elsevier.

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The strong correlation between the d-band center and the bond energies and activation barriers makes it interesting to have an overview of the way the d-band centers change when one metal forms a monolayer on top of (or is alloyed into) the first layer of another metal, these shifts being calculated by using the DFT [16, 18]. All the combinations of metals to the right of and including the Fe group was studied, since these systems contain a number of the metal combinations that could be of interest in catalysis. The main conclusion of Nørskov and co-workers’ work is that when considering variations in the reactivity of a particular metal or group of metals, a single parameter, the d-band center is strongly related to both the stabilities of atoms and molecules on the surface and the energies of transition states for surface processes [16, 18]. For example, the higher the d-band center, the more likely antibonding metal d-adsorbate states are to be pushed above the Fermi level and the more likely the metal d-adsorbate interaction is to become net attractive (that is, the lower the transition energy). When available, the set of both the d-band center of single metals and the d-band centers of alloyed metals is a first tool for understanding and trying to predict the reactivity of the metalbased electrocatalysts reactivity. Accordingly, Nørskov and co-workers [16, 18] suggested a useful table (Table 1). Notice that it has been reported that in most cases the local average of the d-electron energies suffices to describe the electronic effect [16].

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2.2 Segregation A way of changing the reactivity of a given metal is through alloying. The addition of one or more chemical elements to a metallic surface increases the possible bonding geometries of adsorbates and reaction complexes and simultaneously changes the electronic structure of the alloy surface from that of the pure metallic surface alone. The effect of alloying may also be more indirect if, for example, one of the alloy elements segregates to the surface. The chemical composition at the surface of an alloy may then differ from the composition in the bulk; that is, one of the alloy components may enrich the surface by segregation. The tendency of one metal to segregate to the surface of another largely controls the surface composition. This phenomenon, known as surface segregation, is of vital importance in all surface chemistry as it may enhance or suppress desirable and undesirable chemical reactions. In the design of a catalyst, it is therefore of interest to have an overview of the tendency of metals to segregate to the surface. Again, DFT calculations are a useful source of information [16, 19, 22, 23]. Predicting segregation for binary alloys is essential in predicting surface composition of materials since the surface is the active area of the particle. For that, a quite complete data list is available in the literature [16, 22, 24] and is reproduced here (Figure 3). Segregation phenomenon may explain geometric effects. According to Hammer and Nørskov [16], if adsorbates or reaction complexes interact in different geometric arrangements with surface atoms with identical local electronic properties, the differences in chemisorptions bonds and energy barrier heights are ascribed to the differences in the bonding geometry. The simplest measure of the geometric effect is thus the coordination number of the adsorbate with respect to the surface atoms. In addition to segregation predictions, Nørskov and co-workers [23] have also suggested predictions for determining whether two metals of a binary compound will mix in the surface layer or not (data partly reported here in Table 2); knowledge of segregation energy as well as of surface mixing

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energies will be useful as guidance for studies of transition metals in and on other transition metals.

Figure 3.Surface segregation energies of transition metal impurities (solute, horizontal axis) for the closed-packed surfaces of transition metal hosts (vertical axis); Reprinted figure with permission from [22]. Copyright (1999) by the American Physical Society.

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Table 2. Prediction of the mixture of one metal (impurities; columns) segregating from a host (second metal: rows): positive sign to signify that impurity mixes in the surface layer; Adapted in part from reference [23]. Fe

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Fe

Co

Ni

Cu

Mo

Ru

Rh

Pd

Ag

Ir

Pt

Au

+

-

-

+

+

+

+

+

+

+

+

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+

+

+

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+

+

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+

+

+

+

+

+

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+

+

+

-

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+

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-

+

+

+

-

-

-

+

+

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-

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+

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+

+

-

-

Co

-

Ni

+

+

Cu

+

-

+

Mo

+

-

-

-

Ru

+

+

+

-

+

Rh

+

+

+

-

+

-

Pd

+

+

+

+

+

-

-

Ag

-

+

+

+

-

-

+

+

Ir

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+

+

-

+

+

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Pt

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+

+

+

+

+

+

+

-

-

Au

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+

+

+

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-

+

+

+

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+

2.3 Useful Theoretical Tools Results from DFT calculations can be used to analyze and interpret the reactivity of metal catalysts like e.g. methanol oxidation electrocatalysts. In the following sections, the discussion about the variation of the d-band center of the metals, the segregation and the surface mixing of two metals is systematically referred to data given in Tables 1 and 2 and Figure 3. Such operation has already been realized as it is showed below by three different examples. Mavrikakis and co-workers [12] used some results from their DFT-calculations published elsewhere [24, 25] as parameters for building an entirely DFT-based microkinetic model of methanol decomposition on platinum which successfully fitted experimental data collected on a silica-supported platinum catalyst. It was especially showed that the methanol decomposition pathway beginning with C–H bond scission (CH3OH → H2COH → HCOH → CO) emerged as the predominant one and that CO was the most abundant species on the surface under all reaction conditions studied, this being is in agreement with the observation that platinum surfaces are highly poisoned by CO.

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Shimodaira et al. [26] applied DFT calculations to investigate mechanisms occurring on anode electrocatalysts of PEMFCs. They investigated adsorption energies of CO and H2 using Pt-based binary alloys, the second metal being V, Cr, Fe, Co, Ni, Mo, Ru, Rh, Pd, Ag, Sn, W and Au. Most CO durable anode electrocatalysts were searched according to two criteria: that is, CO adsorption energy had to be smaller than that of pure Pt metal; the activation energy for H–H bond fission did not have to be higher than that of pure Pt metal. The first criterion was satisfied by 14 alloy systems and the second criterion by the six systems that were Pt-Co, Pt-Ni, Pt-Ru, Pt-Au, Pt-Fe and Pt-Mo. Except for Pt-Au, Shimodaira et al. [26] showed that their DFT-based screening of Pt-based binary alloys matched experimental findings. It was then suggested that the computational approach including the two criteria may be considered to be useful to search CO durable electrocatalysts. Then, periodic DFT calculations were carried out to explain the surface segregation tendency of one component element over the other for the six alloys. It was found that Pt segregates for Pt-Co, Pt-Ni, Pt-Ru, Pt-Fe and PtMo while it antisegregates for Pt-Au. Note that these observations are rather consistent with the surface segregation energies given by Figure 3. Shimodaira et al. [26] especially showed that Pt-Ru is a CO-tolerant electrocatalyst for e.g. the methanol oxidation (what is really the case according to the articles available in the open literature; see e.g. reference [27]). Furthermore it was suggested that the reaction mechanisms was consistent with the so-called bifunctional mechanism (more details are given below, in section 3.1). Hydrogen is increasingly used in energy production and storage technologies but hydrogen has to be produced. In consequence, hydrogen evolution reaction, which involves proton reduction and concomitant hydrogen evolution, is one of the reactions that has attracted much attention [28], especially in DFT [25, 29], first because of its technological importance, second because it is supposed to be one of the apparently simple electrocatalytic reactions that are suited for model studies, for example, with single crystal surfaces [28]. In other words, theoretical methods can now be used to semi-quantitatively describe the hydrogen evolution reaction [30]. Greeley and Mavrikakis [25] used periodic, self-consistent DFT calculations to develop new near-surface alloys for such processes. Using simple models, the stability of alloys in hydrogen-rich environments were evaluated and the reactivity of hydrogen on the resulting stable surfaces was discussed. For that, Greeley and Mavrikakis [25] conducted a methodic DFT-based screening. First, the surface composition of binary alloys was estimated from the solute/host segregation energy database of Nørskov and co-workers [22]. Second, the effect of hydrogen adsorption on the surface composition was assessed because, in general, adsorbates can draw to the surface components of alloys to which they bind strongly. From that, the hydrogen binding energy on the close-packed surfaces of a wide variety of pure metals and near-surface alloys were assessed. It was demonstrated that near-surface alloys generally bind more weakly than do pure metals. Although weaker hydrogen binding should always imply a higher hydrogen dissociation barrier, near-surface alloys, however, escape from that rule by offering the exciting combination of weak hydrogen binding and easy hydrogen activation. Many near-surface alloys may have favorable hydrogen dissociation kinetics in spite of their relatively weak hydrogen binding. Hence, the authors concluded that stable near-surface alloys could be designed that bind CO weakly. Fortified by the success of the DFT-based screening scheme, Greeley, Nørskov and co-workers [29-31] used these strategies to identify a new bimetallic electrocatalyst for the hydrogen evolution reaction, by theoretically evaluating the activity of over 700 binary alloys. Compared to the article discussed above [25], the reference [29]

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tackled the screening with more theoretical considerations. The Sabatier principle [32] was considered [11, 29]. This principle states that optimal catalytic activity can be achieved on a catalytic surface with intermediates binding energies (of free energies of adsorption) for reactive intermediates. If the intermediates bind too weakly, it is difficult for the surface to activate them, but if they bond to strongly, they will occupy all available surface sites and poison the reaction; accordingly intermediate binding energies permit a compromise between these extremes. The catalytic activity of a material for the hydrogen evolution reaction can be plotted as a function of the hydrogen-metal bond strength, a volcano-shaped form being found. Catalysis for hydrogen evolution reaction is relatively simple, because the catalytic activity can be summarized by a single reactivity descriptor, ΔGH, the free energy of hydrogen adsorption on the catalyst [11, 30, 31]. Basically, for very exothermic hydrogen adsorption (ΔGH < 0), the coverage of adsorbed hydrogen will be high and will, in essence, poison the hydrogen evolution process; for very endothermic adsorption (ΔGH > 0), the model predicts that a high barrier for H* formation from solvated protons will lead to low exchangecurrent densities; for ΔGH about 0, these two regimes are approximately balanced, and a maximum in the exchange-current density is predicted. A good catalyst is characterized by a low activation energy and weak bonding of the intermediates [15]. Therefore, using periodic, self-consistent, DFT calculations, the value of the free energy of hydrogen adsorption on the 736 distinct binary transition-metal surface alloys that can be formed from the 16 metals Fe, Co, Ni, Cu, As, Ru, Rh, Pd, Ag, Cd, Sb, Re, Ir, Pt, Au and Bi, was evaluated. Furthermore the stability of the surface alloys in real electrochemical environments was estimated. Then, 4 simple tests were carried out. First, the free-energy change associated with surface segregation events was estimated. Second, the free-energy change associated with intrastructure transformations such as island formation and surface de-alloying was determined. Third, the free energy of oxygen adsorption, beginning with splitting of liquid water was evaluated (facile oxygen adsorption can lead to surface poisoning and/or oxide formation). Finally, the likelihood that the surface alloys of interest would corrode in acidic environments was estimated. Accordingly, the stability considerations eliminated a large number of alloys from consideration. Only a small fraction were predicted to be both active and stable in acidic hydrogen evolution reaction environments, including, among others, surface alloys of Bi-Pt, Pt-Ru, As-Pt, Sb-Pt, Bi-Rh, Rh-Re, Pt-Re, As-Ru, Ir-Ru, Rh-Ru, Ir-Re and Pt-Rh. Thus, these results demonstrated that stability considerations are essential for finding realistic candidate electrocatalysts for the hydrogen evolution reaction. Furthermore, in order to validate such an analysis, Bi-Pt was synthesized and tested experimentally. This one showed improved hydrogen evolution reaction performance in relation to pure Pt. This experimental observation was in agreement with the computational results. Mavrikakis [11] commented this work by arguing that the work showed that modern electronic-structure theoretical methods, such as DFT, have become affordable and accurate enough for scientists to use them for screening a large number of catalytic materials and for identifying promising alloy compositions, while remarking that theory is not expected to make experiments obsolete in the search for new electrocatalysts; it is expected, however, to become increasingly important in guiding efforts in experimental synthesis in the right direction. Hence the DFT calculations-based results are used in section 3 as theoretical tools for attempting to explain the reactivity of metal electrocatalysts envisaged as anode electrocatalysts in some DLFCs.

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3. METAL-BASED ELECTROCATALYSTS FOR DLFCS 3.1 Direct Methanol Fuel Cell

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Methanol is a potential fuel of DLFC because its oxidation reaction is a process involving the transfer of six electrons per molecule to the electrode: (1) Anode reaction: CH3OH + H2O Æ CO2 + 6 H+ + 6 eCathode reaction: 3/2 O2 + 6 H+ + 6 e- Æ 3 H2O

(2)

Overall reaction: CH3OH + 3/2 O2 Æ CO2 + 2 H2O

(3)

The standard cell voltage is 1.21 V at 25 °C. The overall reaction (3) is the ideal one where the reaction products are only carbon dioxide and water. In terms of greenhouse gas (i.e. CO2) emissions, the methanol oxidation reaction is less polluting than the gasoline combustion [7]. Moreover, when methanol is generated from renewable raw materials (manure, sewage, landfills, etc.), the CO2 emissions does not contribute to any additional greenhouse effect [33] (if the CO2 emissions related to the methanol production processes and the transport are not considered). The DMFC is certainly the most promising DLFC [34] but in order to be competitive it must be reasonably cheap and capable of delivering high power densities. However, with respect to this latter point there are few challenging problems, one of these being the efficiency of the anode electrocatalysts [27, 34, 35]. Intensely investigated, lots of papers dealt with this technology and some review papers have been published. For example, Aricò et al. [27] described developments in both the fundamental and technological aspects of DMFC, focusing e.g. the anodic electrocatalysts. More recently, Demirbas [34] briefly reviewed the main advantages of the DMFC and emphasized on the Pt-based electrodes that has attracted considerable attention in the methanol oxidation. Aricò et al. [27] largely discussed the anodic oxidation of methanol and in that way, commented the state-of-the-art of the anode electrocatalysts. Pt is a very efficient metal for catalyzing the methanol oxidation (Figures 4 and 5) [36] but unfortunately various reaction intermediates (that is, CO-like species) may form. These species are detrimental to Pt because, irreversibly adsorbing on the metal surface, they severely poison it. As a consequence, the DMFC performance deteriorates. Therefore new electrocatalysts have been envisaged and such an action has two goals: to inhibit the poisoning due to strong adsorption of CO-like species and to significantly increase the oxidation kinetics. Aricò et al. [27] briefly reviewed the electrocatalysts alternative to Pt and, from that review, some general observations emerge. From a general point of view, the methanol oxidation requires the presence of Pt-based electrocatalyst but Pt strongly chemisorbs CO. The efficiency and selectivity of electrocatalytic processes may be substantially improved by replacing monometallic with bimetallic electrocatalysts [37]. Partial substitution of Pt with other transition metals like e.g. Pd or Ni, or alloying it with other elements is a solution. It has been remarked that the most successful performances have been achieved through the alloying route. The alloying of Ru with Pt has given rise to electrocatalysts which strongly promote

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the oxidation of both methanol and CO. For example, Liu et al. [38] showed that the binary Pt-Ru electrocatalyst was more active than the Pt-only electrocatalyst and that the former was less receptive to methanolic residue deactivation. Fujiwara et al. [36] compared Pt, Ru, Rh, Ir, Pd, Au and Pt-Ru (Figures 4 and 5). The oxidation current on these metal electrodes obtained in cyclic voltammograms and the maximum power densities of DMFCs constructed with these metals as anode electrocatalysts showed that Pt-Ru was the best methanol oxidation electrocatalyst, followed by Pt; Ir was quite inactive and the other metals were inactive. Aricò et al. [27] discussed the experimental methods and the electrode kinetics for introducing a discussion about the improved resistance to poisoning of the bimetallic alloys. Reaction mechanisms were also suggested. For Pt-Ru alloys, Pt sites are involved in both the methanol dehydrogenation and in the strong chemisorption of methanol residues (like e.g. CO) while Ru is involved in water discharging with the formation of Ru-OH groups at the electrocatalyst surface; these groups react with neighboring methanolic residues adsorbed on Pt to give carbon dioxide (4) through a bifunctional mechanism (each metal component promotes different elementary reaction steps [8, 37, 38]):

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Ru-OH + Pt-CO Æ Ru + Pt + CO2 + H+ + e-

Figure 4. Oxidation current of fuel compounds on various electrodes obtained in multiple-cyclic voltammograms at 25 °C; Reprinted from [36], Copyright (2007), with permission from Elsevier.

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(4)

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Figure 5. Maximum power densities of DLFCs with various fuel compounds and anode electrocatalysts at room temperature and atmospheric pressure; Reprinted from [36], Copyright (2007), with permission from Elsevier.

According to Aricò et al. [27], the promotional effect of Ru would not be only due to the bifunctional effect but there would be the contribution of a ligand effect. In other words, an electronic effect resulting from interactions between the two metals could improve reactivity [37]. However, Strasser [39] showed that the most active Pt-Ru binary alloy was the Pt88-Ru12 composition and the less active was Pt52-Ru48. The author remarked then that these results seemed to contradict the bifunctional rational. It was suggested that methanol requires an ensemble of three adjacent Pt surface atoms for its adsorption and subsequent hydrogen stripping toward CO; once the CO has formed on an ensemble of three Pt atoms, it needs an adjacent Ru atom providing OH species to complete the reaction. Strasser’s results [39] are consistent with earlier studies on Pt-Ru electrocatalysts [40, 41]. Interestingly, whatever the preparation route, Pt-Ru has always been the best performing Pt-based alloy. From that, PtRu has been an object of near exclusive attention and in order to even improve its performances investigations the optimization of the preparation conditions has been focused (see for example reference [42]). Note that the present article does not review the electrocatalyst preparation methods because of a huge body of literature and some existing reviews devoted on these [9, 43-46]. Pt-Ru may be either an alloy or a binary system i.e. a mixed monometallic systems of the same composition. Recently, a Ru@Pt core-shell nanoparticle catalyst that is distinctly different from the previous Pt-Ru catalysts was reported in the preferential oxidation of carbon monoxide [47]. Although the methanol was not envisaged in that study, it is important to report the main observation of that paper since it discusses carbon monoxide (poison in the

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DMFC) as well as its oxidation (the reaction at the origin of CO-like species removal from the metal surface). It was showed that the Ru@Pt nanoparticles are far more active for CO oxidation than alloy and monometallic nanoparticle catalysts. For this catalytic material, the Ru metal was confined and kinetically trapped inside a Pt shell and then the bifunctional mechanism could not be implicated because CO oxidation necessarily occurred entirely on the Pt surface sites. No Ru was exposed on the surface of the nanoparticle. In order to explain the enhanced CO oxidation of Ru@Pt, DFT modeling were used and it was showed that the enhanced reactivity was achieved through modification of the electronic structure of the Pt surface by the presence of subsurface Ru. This modification significantly destabilized CO on Pt, leading to a lower CO saturation coverage (i.e. a decrease in the interaction strength of CO adsorbates on the Pt surface atoms [48]), thereby providing more adsorbate free active sites where reactants could be activated. At the same time, this electronic modification greatly accelerated the CO oxidation reaction through a substantial destabilization of the adsorbed reactive intermediates. The interaction of the Pt-monolayer with the Ru-support atoms caused a downshift in the d-band centre of surface Pt atoms. Downshifting the d-band centre of a surface was shown to decrease the interaction strength of the surface with various adsorbates [18, 48]. Therefore, this study as well as the studies reported above show that at high CO concentration environments Pt-Ru or Ru@Pt will have significantly more CO-free sites to carry out catalytic reactions than the CO-saturated Pt surface. Demirci [14] also tackled the catalytic efficiency of Pt-Ru. The Pt-Ru reactivity is especially discussed on the basis of density functional calculations of adsorption and reactions at metal surfaces. Precisely, Demirci [14] took into consideration the following suggestions: alloying may increase the possible geometries of the adsorbates and reaction complexes and may change the electronic structures of the alloy in relation to those of the pure metallic surfaces of the alloy constituents; the surface segregation may control the surface composition of the alloy. It has especially been showed that alloying Pt with Ru implies a change in the electronic structure of Pt; that is, when Pt and Ru are alloyed, the adsorption of adsorbates is weaker on the Pt sites and stronger on the Ru sites. Because of such electronic effects, the Pt poisoning by the CO-like species would be decreased because these species would adsorb more weakly and Ru would provide the necessary OH species, which would be permanently available on the Ru surface sites, more strongly adsorbing the OH species. Aricò et al. [27] had also reported this electronic effect. In fact, it was suggested that alloying Pt with Ru produced an increase in Pt d-band vacancies and this might modify the adsorption energy of methanolic residues on Pt. Accordingly, the reaction rate would not be only dictated by the bifunctional mechanism but also by the electronic effects on account of the interaction between Pt and Ru. According to the DFT-based segregation predictions, Pt is expected to segregate while Ru antisegregates. Besides, Pt and Ru mix in the surface layer what suggests that Pt and Ru coexist together at the metallic surface although surface concentration of Pt is higher than that of Ru [23]. These geometric predictions are in agreement with the bifunctional mechanism taking place over Pt-Ru. With respect to CO poisoning noticed with Pt electrocatalysts, Greeley and Mavrikakis [13] brought to that discussion an approach more theoretical that is based on DFT calculations. Since the anode must be resistant to CO poisoning, this property could be realized, for example, by finding alloys that bind CO weakly. As that was the case with hydrogen, the binding of CO to the near-surface alloys is generally weaker than is the CO binding to the corresponding host metals. However, the alloys must not be suffering from CO-induced surface segregation. Indeed CO will, therefore,

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possess a larger thermodynamic driving force for adsorbate-induced segregation. Greeley and Mavrikakis [13] suggested that one of the promising binary alloys, resistant to CO poisoning, is Pt-Ru. CO-induced segregation should not take place with Pt-Ru and thus this bimetallic combination should be stable. Furthermore, the CO binding energies of some bimetallic alloys were assessed and it was showed that the binding over Pt of Pt-Ru is lower than that over pure Pt. Note that for this Pt-Ru alloy the H-induced segregation has not been envisaged [25]. All of these assumptions would let one think that Pt-Ru is a stable alloy and even in the operating conditions of the methanol oxidation. However recent durability tests showed Pt and Ru from Pt-Ru separately moved through the DMFC membrane to the cathode [49]. This asks then a question about the stability of Pt-Ru: may adsorbate-induced segregation be at the origin of such deterioration? Further studies are required for getting answers to such interrogations. Note that besides Pt-Ru several Pt-based bimetallic alloys (e.g. Pt-Co, Pt-Ni, Pt-Rh, or PtAu) were discussed in similar terms [14] and investigations about the combination Pt-Sn were reviewed [27]. Nevertheless Pt-Ru is undeniably the best bimetallic alloy. Globally, the DFTbased calculations provide satisfying interpretations of the better reactivity and efficiency of Pt-Ru. A way of improving its reactivity and stability is to add a third metal. Indeed surface or subsurface atoms of additional metals result in modifications of the electronic or structural characteristics of the single or bimetallic alloy surface [50]. Pt-Ru-based trimetallic alloys were interestingly tackled by Aricò et al. [27] and Demirci [14]. According to the former authors, the third metal brings a third functionality but the authors do not give any further details about this functionality. Lamy et al. [10] studied few Pt-Ru-based trimetallic alloys; indeed Au, Co, Cu, Fe, Mo, Ni, Sn and W were added to Pt-Ru. Globally the addition of a third element showed promising results. The best promoter was found to be Mo, followed by W, Co, Fe, Ni and Cu. The presence of Sn or Au was detrimental since the trimetallic material showed lower performances than the bimetallic Pt-Ru. Demirci [14] attempted to discuss the trimetallic electrocatalysts reactivity by using theoretical concepts and suggested that the improvement of the Pt-Ru alloys would require a third metal with which Pt and Ru segregate and their adsorption strengths weaken. In other words, Pt-Ru-based trimetallic alloys should have Pt-and-Ru-enriched surfaces, where the Pt sites would be less sensitive to the CO-like species adsorption. Recently Strasser [39] reported the combinatorial and highthroughput optimization of improved ternary Pt-Ru-Co alloy electrocatalysts. In addition to pure Pt and a number of Pt-Ru binary alloys, 64 ternary alloy electrocatalysts were synthesized to optimize the Pt20-Co60-Ru20 that had showed promising electrocatalytic performances in a previous study [51]. First, the methanol oxidation activity was assessed and a composition-activity map was proposed: several electrocatalysts neighboring Pt20-Co60-Ru20 alloy did exhibit a significantly higher activity than pure Pt. Second, the corrosion stability of the synthesized ternary materials was studied and a composition-stability map was proposed. Finally, a consensus map by combining the previous two maps was showed in order to conveniently facilitate the down selection of suitable preferred lead alloy candidates (Figure 6). The Pt18-Co62-Ru20 alloy exhibited the highest methanol oxidation activity at acceptable stability inside the activity-stability consensus region. At the knowledge of the present author, Strasser’s study [39] is the most succeeded work about trimetallic Pt-based electrocatalysts and the successful utilization of combinatorial optimization is an evidence of the power of such tool.

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Pt

Co

Pt Pt16Co62Ru20

Co

Ru

Ru

Figure 6. Consensus map of the Pt-Co-Ru focus library superimposing activity and stability maps. The dashed circled region indicates the compositional region of favorable activity, while the gray shaded area represents the compositional region of severe corrosion and instability. The region delineated by the solid black line represents the preferred consensus region. The alloy Pt18-Co62-Ru20 is selected as the lead electrocatalyst on the basis of a balance of activity and stability considerations; Reprinted figure with permission from [39]. Copyright (2008) by the American Chemical Society.

3.2 Direct Ethanol Fuel Cell Ethanol, safer than methanol, is also a potential fuel of the PEMFC [7]. Basically the direct ethanol fuel cell (DEFC) is based on the oxidation of ethanol, the overall reaction being as follows:

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C2H5OH + 3 O2→ 2 CO2 + 3 H2O

(5)

DEFC is facing some issues: e.g. slow kinetics and electrocatalyst poisoning by CO-like species at low temperatures [52-55]. The ethanol oxidation is more complicated than the methanol oxidation because the breaking of the C–C bond to obtain its complete oxidation is necessary. It is then expected from the anode electrocatalyst of the DEFC cleaving of the C–C bond while oxidizing the CO-like intermediates [56]. Like for the DMFC, the most common anode electrocatalyst for the DEFC is Pt but, as discussed above, Pt is too sensitive to the CO-like species and therefore to poisoning [10, 56]. Multifunctional electrocatalysts were then envisaged [56, 57]. Pt-based binary compounds were first assessed [54, 56]. In fact, whatever the direct alcohol-feed fuel cell is, either the DMFC or the DEFC, there is no fundamental difference in the properties sought in the Ptbased alloys: the main objective is to modify the electrode surface in such a way to increase its coverage in adsorbed OH coming from water, the OH species being necessary to oxidize the CO-like species coming from the dissociation of ethanol [14, 56]. Among the investigated materials, Pt-Sn showed the best results [10, 53-55, 58, 59]. Pt-Sn enhanced the CO oxidation to a larger extent than did Pt-Ru and (as expected) Pt [57]. Despite better results with Pt-Sn, some investigations focused Pt-Ru (Figures 4 and 5) [36]. It is globally showed that the Ru atoms act mainly by providing OH species for the oxidation of ethanol intermediates [60].

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Other binary Pt-M (with M as W, Pd, Rh, Re, Mo, Ti, Ce) electrocatalysts were investigated but these electrocatalysts showed an ethanol oxidation reactivity lower than that of both Pt-Sn and Pt-Ru [56] and sometimes than that of Pt. For example, Bergamaski et al. [61] showed that the addition of Rh to the Pt electrocatalyst improved the efficiency for the oxidation of ethanol to CO2 but this better efficiency was not followed by higher faradaic current densities. The authors suggested then that the reduced faradaic current for ethanol oxidation may be caused by the decrease of the Pt content in the bimetallic electrocatalyst or it could be caused by some Rh segregation to the nanoparticle surface. From the prediction of the segregation occurrence, the latter suggestion may be unlikely since Rh is expected to antisegregate when alloyed to Pt. Some adsorbate-segregation might besides occur although H- or CO-induced segregation is not reported [13, 25]. Bergamaski et al. [61] attributed to the better activity for CO2 formation of the Pt-Rh electrocatalyst compared to that of pure Pt the so-called ligand effect, that is, an electronic effect by which Rh modifies the electronic properties of Pt in such a way that the Pt-adsorbates interaction is weakened. This explanation is the same than that suggested for the Pt-Ru and Pt-Sn bimetallic electrocatalysts used in DMFC and DEFC respectively. Lima and Gonzalez [60] also investigated Pt-Rh as anode electrocatalyst of the DEFC. Interestingly they supported their study with the Nørskov and co-workers’ DFT-based calculations that are also discussed here. It was showed a pronounced electronic modification of the Pt 5 d-band when Pt was alloyed with Rh. This led, eventually, to decreased adsorption strength of adsorbates on the Pt atoms. Note that the Pt d-band center is expected to shift up when it is alloyed to Rh and therefore the Lima and Gonzalez’s observations agree the DFT-based predictions. These authors suggested that the Rh atoms had an important role on the C–C bond dissociation. Sn-based bimetallic electrocatalysts were besides investigated. It was showed that the overall performance of Ir3-Sn was comparable to that of Pt3-Sn, making it a promising alternative choice of anode electrocatalyst for DEFCs [62]. With respect to Pt-Sn, Sn has the same role than Ru has in Pt-Ru used in the DMFC; that is, Sn supplies OH species for the oxidative removal of CO-like species strongly adsorbed on adjacent Pt sites [56, 63]. Pt-Sn improves the removal of adsorbed CO species [55] and this has been attributed to changes in the geometric and electronic nature of the surface [64]. For example, Sn may decrease the size of surface Pt ensembles [65, 66] and CO-like species may adsorb more weakly on Pt-Sn electrocatalysts compared to Pt electrocatalysts [67]. Indeed, with Pt-Sn, the poisoning by CO is greatly reduced and a significant enhancement of the electrode activity takes place. However, the oxidation of ethanol is incomplete and C2 products form [57]. Two main routes would exist for the overall mechanism of the ethanol oxidation over Pt-Sn: first, formation of adsorbed CO; second, formation of C2-species like e.g. acetyl species, which lead to the formations of acetaldehyde and acetic acid as by-products [54]. Demirci [14] attempted to explain the Pt-Sn reactivity in terms of electronic and geometric effects by using results from DFT-based calculations. However that attempt failed because of a lack in information regarding Sn in the databases that were considered. Nevertheless, from thermodynamical data, it was suggested that Sn would segregate while Pt would antisegregate. Recently, Kim et al. [68] provided highlighting about the geometric and electronic effects for the Pt-Sn electrocatalyst (studied for the oxidation of ethanol as well as of both methanol and propan-1ol). For the aspect of geometric structure modification on Pt-Sn, XRD and EXAFS analyses indicated that the geometric environment was changed with Sn addition to Pt by forming solid solution of Pt-Sn alloy phase, accompanying an expansion of lattice parameter. This

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elongation of bonding structure may affect catalytic reaction pathways that require specific geometric arrangements of the surface atoms, thus leading to change of catalytic properties. It was besides showed the preferential Sn enrichment on the surface with Sn addition (Figure 7). This last observation is in agreement with Demirci’s predictions [14] as well as the explanations given by some of the authors cited above in this section. With respect to the electronic effect with Pt-Sn, from XPS and XANES results, Kim et al. [68] suggested the charge transfer from Sn atoms to neighboring Pt atoms. In other words, the Sn addition modifies the electronic environment of Pt-Sn electrocatalysts by influencing the electron affinities of the surface Pt atoms and produces surface oxygenated species which can act as an oxidant source, thereby increasing stability through efficient oxidation of surface poisoning species. In terms of d-band center, one might suggest that when Sn is alloyed with Pt the dband center of the latter metal shifts up. With respect to Ir3-Sn, it was interestingly showed a catalytic performance that was comparable to that of Pt3-Sn [62]. It may be interesting to consider the segregation occurrence for Ir3-Sn. By taking into account the heats of sublimation of Ir and Sn (156 and 60 kcal mol-1, at 25 °C, respectively [69]), Sn should segregate (the metal with the lowest heat of sublimation is expected to segregate). It is besides to note that Ir and Pt have quite similar d-band centers (-2.11 and -2.25 eV, respectively). Hence, Sn addition to either Ir or Pt might have comparable electronic and geometric effects on Ir and Pt reactivity. Unfortunately, the lack of theoretical data with respect to Sn does not enable to discuss more deeply such likely effects. Whatever it may be, it is undeniable, from experiments, that Sn improves both reactivity and stability of Pt (and of Ir). For the DEFC, the objective is to find an electrocatalyst that is active towards dehydrogenation, C–O and C– C bond cleavages during the ethanol oxidation and moreover that is resistant to CO poisoning. Hence an attempt in order to identify potential bimetallic alloys from theoretical databases was started on, but no bimetallic combinations fitted the requirements in terms of electronic and geometric effects [14]. It was then concluded that the best choice in the development of new electrocatalysts for the DEFC would be the addition of a third element to either Pt-Sn or Pt-Ru. From the present discussion, one may also suggest Ir-Sn for preparing trimetallic alloys.

Figure 7.Sn/Pt atomic ratios on the surface calculated by XPS analysis with increasing Sn/Pt atomic ratios obtained by ICP; Reprinted from [68], Copyright (2008), with permission from Elsevier.

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Regarding the trimetallic alloys, Demirci [14] attempted to suggest potential candidates to assess but unfortunately no proposition of Pt-Sn-based trimetallic materials was suggested because of the absence of theoretical data about Sn [16, 22, 24]. With respect to Pt-Ru, it was suggested to avoid Pd, Ag and Au (because of Pt dilution in their presence) while envisaging perhaps V, Cr, Mo, Tc, W, Re and Os (because of Ru segregation). Antolini [56] suggested that it is crucial to add a third element to both Pt-Sn and Pt-Ru electrocatalysts to improve their activity. The ternary Pt-Ru- and Pt-Sn-based electrocatalysts were reviewed [56]. All ternary Pt-Ru-based electrocatalysts better performed than Pt-Ru. Pt-Ru-Ni was investigated and its performance for ethanol oxidation was better than that of the Pt-Ru electrocatalysts due to, according to the authors, the promoting effect of Ni [56, 70]. Moreover the COtolerance performance of the trimetallic electrocatalysts was better than that of Pt-Ru [71]. Ribadeneira and Hoyos [70] evaluated 10 different binary and ternary electrocatalysts made of mixture of Pt, Ru, Sn, or Ni and it was demonstrated that Ni addition improved the catalytic activity of both Pt-Ru and Pt-Sn. Furthermore, trimetallic catalytic mixtures performed better than bimetallic electrocatalysts. The following ranking was proposed (partially reported here): Pt75-Ru15-Ni10 > Pt75-Sn15-Ni10 > Pt75-Ru15-Sn10 > Pt75-Sn15-Ru10 > Pt75-Sn25 > Pt75-Ru25. Ni segregates in the presence of Ru while antisegregating when alloyed to Pt. Meanwhile Ru antisegregates in the presence of Pt. These segregations suggest that the surface alloy may basically be constituted of three layers from the surface, each layer being enriched with a metal: Pt layer at the top surface, Ni and then Ru sublayers. Furthermore, Ni mixes with Pt while not mixing with Ru. With respect to the d-band centers, the Pt d-band center when alloyed with Ni shifts downer than that when alloyed with Ru. This suggests that Pt should binder much more weakly CO-like species. From these observations, one may suggest that Pt-Ru-Ni is more reactive and stable because of Pt surface sites more weakly adsorbing the reaction intermediates and Ni sites efficient in dissociating water molecules and then in providing OH species. However it is generally reported that Pt-Ru is less reactive than Pt-Sn. The addition of a third metal (e.g. Ir [72], W [73], etc) to Pt-Sn improves the reactivity of the bimetallic electrocatalyst. The best combination is Pt-Sn-Ru but the content of Ru has appeared to be critical [74]. It was suggested the following ranking for Pt-based bimetallic and trimetallic electrocatalysts in the ethanol oxidation (ranking according to the power density that were achieved): Pt, Pt-Pd < Pt-W < Pt-Ru < Pt-Ru-Mo < Pt-Ru-W < Pt-Ru-Sn < Pt-Sn [75]. With respect to Pt-Ru-Sn, if one regards the heat of sublimation of the metals (155, 116 and 60 kcal mol-1 at 25 °C for Ru, Pt and Sn [69]), Ru enriches the bulk while Sn segregates and Pt dilutes at the surface. Though no datum about the d-band center of Sn is available, one may likely suggest that Ru in the bulk may modify Sn reactivity by improving its ability to provide more OH species. Antolini [56, 74] largely discussed the positive effects of both Ru and Sn on Pt reactivity; basically the role of both Ru and Sn is to supply OH species for the oxidative removal of CO-like species strongly adsorbed on adjacent Pt sites. According to Ribeiro et al. [73], the presences of Sn and W favor the activation of interfacial water molecules, which are necessary for the removal of irreversibly adsorbed species such as linearly adsorbed CO. Colmati et al. [76] modified Pt-Sn by adding Rh. The ternary Pt-Sn-Rh alloy electrocatalyst possessed a high activity for ethanol oxidation, higher than that of Pt-Sn. The enhanced activity was ascribed to changes in the geometric and electronic structures due to the formation of a ternary alloy and to lower particle sizes. Geometric and electronic effects may explain reactivity changes in trimetallic alloys in comparison to that of bimetallic

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electrocatalysts but for characterizing such effects DFT-based combinatorial optimization is certainly the most efficient tool.

3.3 Direct Propanol Fuel Cell

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Alcohols alternative to methanol and ethanol can only be higher molecules. Both propan1-ol and propan-2-ol have been evaluated as fuels for the PEMFC [7, 77]. As that could be expected, the oxidation of higher alcohols always produced some amounts of adsorbed CO [10]. Kim et al. [68] evaluated the electrocatalytic activity of Pt-Sn towards the oxidation of methanol, ethanol and propan-1-ol (Figure 8). Wang et al. [78] evaluated ethanol, propan-1-ol and propan-2-ol by using Pt–Ru and Pt electrocatalysts. For propan-1-ol, propanal and CO2 were mainly produced. For propan-2-ol, the main by-products were propan-2-one and negligible amounts of CO2. Both C3-alcohols did not appear to be suitable fuels because of their low electrochemical activity. In fact, the main difficulty with these fuels is the cleavage of one C–O and two C–C bonds. Alcohols undergo adsorption and dehydrogenation steps on the active Pt sites to give adsorbed reaction intermediates before the cleavage of C–C and C– O bond, while the dissociative adsorption of water is believed to occur on the Ru or Sn sites to form surface oxygen-containing species, in which the latter are expected to be able to facilitate the oxidative removal of the intermediate residues [68]. However, Kim et al. [68] emphasized that the electrochemical oxidations of the alcohols that possess several carbon (2, 3 or 4) atoms are still under discussion because of much complicated reaction mechanisms. According to the knowledge of the present author, literature that is devoted to these alcohols is still at its beginning and therefore very few papers are available. Furthermore, propan-1-ol and propan-2-ol seem to be much less promising than both methanol and ethanol as liquid fuels of DLFCs.

Figure 8. The electrocatalytic activity as a function of Sn/Pt ratios for methanol, ethanol and propan-1ol; Reprinted from [68], Copyright (2008), with permission from Elsevier.

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The first investigations used Pt and Pt-Ru as anodic electrocatalyst because of their known catalytic abilities in both DMFC and DEFC. For example, Fujiwara et al. [36] compared Pt, Ru, Rh, Ir, Pd, Au and Pt-Ru (Figures 4 and 5). The propan-2-ol oxidation current on these metal electrodes obtained in cyclic voltammograms and the maximum power densities of the propan-2-ol-powered fuel cell constructed with these metals as anode electrocatalysts showed that Pt-Ru was the best oxidation electrocatalyst, followed by Pt; Ir was quite inactive and the other metals were completely inactive. It is quite evident that the propanol-powered fuel cell will have the same problems than those reported in DMFC and DEFC. Pt-Sn electrocatalysts with various Sn contents were studied for the oxidation of propan-1-ol [68]. The electrocatalysts with low Sn content showed low activities and but they were increased as the Sn content was increased up to the optimum Sn contents (Figure 8), i.e. 40 wt.% for propan-1-ol (it was of 25 wt.% for methanol and 33 wt.% for ethanol). Further Sn increases, however, caused the current densities to decline. Kim et al. [68] rationalized these features as follows: at low Sn content, there may be insufficient Sn sites to provide surface oxygen-derivatives capable of helping oxidize the adsorbed reaction intermediates; around optimum Sn content, a suitable modification might be satisfied in terms of the electronic and geometric structures for the alcohol oxidation; too higher Sn content may seriously inhibit the dissociative chemisorption of alcohol due to too much diminished Pt ensembles size and a substantial reduction of available surface Pt active sites by larger amounts of Sn species. Consequently, a compromising balance between the structural and electronic properties with variation of Sn content may be decisive factors to control their catalytic properties and electroactivities over the different alcohols. In fact, the roles of Pt and Sn are considered as being still the same whatever the alcohols (methanol, ethanol, propan-1-ol) are, besides knowing that these roles are commonly admitted for most of the authors investigating Pt-Sn electrocatalysts (see section 3.2 and the references given therein). Kim et al. [68] summarized that as follows. As a consequence of the increased Sn content, the modification of structural and electronic features were observed in that the increased lattice parameter might activate C– C cleavage and the increased amount of surface oxygen containing species facilitate to oxidize CO-like species structurally, and the charge transfer from Sn to Pt might cause a prevention or reduction of catalytic poisoning by partially filling of the Pt d-band vacancies electronically. Recently, Xu et al. [79] and Liu et al. [80] showed that Pd is a good electrocatalyst for propan-2-ol oxidation in acidic medium and alkaline medium respectively. Liu et al. [80] showed besides that Pd is also a good electrocatalyst for propan-1-ol oxidation in alkaline medium. Xu et al. [79] reported the electrocatalytic activity of both Pd and Pd-Au materials towards the oxidation of propan-2ol. The addition of Au to Pd significantly promoted the electrocatalytic activity and poisoning tolerance of Pd electrocatalysts. Furthermore, both Pdbased electrocatalysts showed better activity and stability than that of a commercial Pt electrocatalyst. With respect to Pd-Au, Xu et al. [79] recognized that the promoting mechanism of Au in Pd-Au electrocatalysts is not clear but despite that they still suggested that Pd acts as primary active sites for the catalyzing of the dehydrogenation of propan-2-ol and CO-like intermediate species can be oxidized on the Au surface. To emphasize this hypothesis, it was remarked that Au is a good catalyst for CO oxidation [81]. In fact, this mechanism is none other than the bifunctional mechanism that is generally accepted for e.g. Pt-Ru in the methanol oxidation. From the predictions about the d-band center variations and the segregation, the d-band center of Pd shifts up when alloyed to Au and Pd antisegregates

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while surface Au and Pd mix. In other words, the surface is enriched with Au while the Pd sites are diluted in the Au matrix (few Pd atoms surrounded of several Au atoms) and the adsorption of species over Pd is weakened. These electronic and geometric effects rather agree the occurrence of the mechanism above suggested. Compared with the DMFC and the DEFC, very few papers reported work about the DPFC. The early results showed that propanol would not be a suitable fuel of DLFC. It is likely that the best bimetallic materials would be Pt-Ru or Pt-Sn, these materials having showed their respective abilities in DMFC and DEFC. In fact, the potential electrocatalyst should be able to cleave two C–C bonds in addition to the C–O bond. This constraint requires thus a highly reactive metal in the C–C bond cleavage. Note that Pt has already showed the best performances. Investigations are in progress and these should next be followed while especially focusing the bi- and tri-metallic electrocatalysts.

3.4 Direct Acid Formic Fuel Cell Another interesting, promising liquid fuel is acid formic [7]. Recently, the direct formic acid fuel cell (DFAFC) has been the object of a review paper, mainly focusing on the anodic electrocatalysts for the oxidation of formic acid [82]. The overall reaction of the DFAFC is as follow: HCOOH + ½ O2 → CO2 + H2O

(6)

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The mechanism of the formic acid oxidation on Pt follows a dual pathway: that is, a dehydrogenation reaction (7) or a dehydration reaction (8) [83, 84]; the dehydration path has to be avoided because of CO formation, detrimental to Pt electrocatalysts: HCOOH → CO2 + 2 H+ + 2e−

(7)

HCOOH → COads + H2O → CO2 + 2 H+ + 2 e−

(8)

In addition to this dual pathway, Chen et al. [85, 86] have more recently proposed a new scheme for the mechanism of formic acid oxidation on Pt, so-called triple pathway involving the bridge-bonded formate formation and subsequent oxidation of formate to carbon dioxide. As it is generally the case for most of new liquid fuels [14], the metals known for their high electrocatalytic abilities were assessed in the oxidation of formic acid. Indeed Pt- and Pd-based electrocatalysts are commonly used in DFAFC, with use of Pd becoming more prevalent [82]. Yu and Pickup [82] reviewed these catalytic materials. Surprisingly, Pd overpassed Pt. For instance, Fujiwara et al. [36] compared Pt, Ru, Rh, Ir, Pd, Au and Pt-Ru (Figures 4 and 5). The formic acid oxidation current on these metal electrodes obtained in cyclic voltammograms and the maximum power densities of DFAFCs constructed with these metals as anode electrocatalysts showed that Pd was the best oxidation electrocatalyst, followed by Pt-Ru and Pt. Thus the efforts brought to Pd and to Pd-based bimetallic combinations. For instance, Pt-Pd electrocatalyst was investigated [87, 88]. It was more efficient than both Pt and Pt-Ru [88]. Pd not only exhibited activity for formic acid oxidation but also provided a synergetic effect to Pt, which oriented Pt-Pd in favor of formic acid

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oxidation via the direct path [87]. Demirci [14] attempted to understand the reactivity of PtPd in terms of electronic and geometric effects. It was reported that Pd and Pt do not segregate (they mix together in the surface layer) and that the surface of Pt-Pd is occupied by both metals in the proportions of the alloy. Note that there is almost no variation of the d-band center of both metals when one is alloyed with the other. It was suggested that Pd do likely not act as an OH source because it is not required for the direct oxidation of formic acid, but Pd likely influences the surface reactions by geometric effect. Pd would modify the Pt environment and so the Pt reactivity. Pd having showed its promising abilities in the formic acid oxidation, Pd-based bimetallic combinations alternative to Pd-Pt were envisaged. Accordingly, Pd-Co [89], Pd-Ir [90] and Pd-Au [91] were investigated. With respect to Pd-Co-based alloys, Pd0.9-Co showed better performance towards the anodic oxidation of formic acid than Pd and Pt [89]. Theoretically Pd segregates while Co antisegregates and then Co enriches the particles sublayers. Moreover the d-band center of Pd shifts up what implies that Pd binds more strongly formic acid than does pure Pd. Wang et al. [89] added a third metal, Ir, to the Pd-Co electrocatalyst; Pd4-Co1.9-Ir0.1 showed better performance towards the anodic oxidation of formic acid than Pd, Pt and Pd0.9-Co. Such enhanced reactivity was attributed to the Ir and Co doping. It was reported that Ir can promote the oxidation of formic acid at Pd because Ir can decrease the adsorption strength of CO. However when the Ir content was too high the bimetallic activity was decreased, what was attributed to the catalytic inactivity of Ir [90]. With respect to Pd-Au [91], it generally exhibited a much lower activity than Pd-Pt. This may be explained by the noble character of Au that then does not catalyze formic acid oxidation as well as Pt. Baldauf et al. [91] emphasized the remarkable resistivity of Pd overlayers against poisoning due to CO-like species. Studies focusing Pt were also led. In their review paper, Yu and Pickup [82] noticed that a series of Pt-based bimetallic systems including Pt-Ru, Pt-Bi, Pt-Pd, Pt-Au, Pt-Pb, Pt-Ir have been reported as effective electrocatalysts for DFAFCs. Choi et al. [92] compared Pt-Au and Pt-Ru. Pt-Au revealed better performances. It was suggested that the enhancement of formic acid oxidation is likely due to the so-called third-body effect, which meant that the addition of a second element (third body) to Pt reduces the number of adsorption sites for CO due to geometrical hindrance and therefore the surface is poisoned by CO to a lesser extent than the surface of pure Pt. Lee et al. [93] also showed that the Pt-Au has a higher activity than those of both Pt and Pt-Ru. Such better results for Pt-Au were also reported by Wang et al. [94] and this increased reactivity was explained by the fact that the formation of the poisoning intermediate CO is suppressed. In the presence of Au, Pt strongly antisegregates and both metals do not mix in the surface layer. The Pt d-band center shifts up, while the Au d-band center remains almost constant. These observations, partly reported by Demirci [14], are in agreement with the explanation provided by Choi et al. [92]. Note that this analysis in terms of d-band shift is similar to that done for Pd-Co in the previous paragraph. The explanation provided by Choi et al. [92] may also be applied to Pd-Co [89]. In addition to Pt-Ru and PtPd, Bi- and Pb-modified electrodes were largely tested [95-102]. These Pt-Bi and Pt-Pb electrocatalysts demonstrated higher activity than the pure Pt surface and this electrocatalytic enhancement was mainly ascribed to reduced poisoning formation because of the electronic effect of the second metal, i.e. Bi or Pb. Such electronic effects are very likely but unfortunately data relative to both the d-band center and the segregation are not available for both Bi and Pb. Note that Nørskov and co-workers [29] selected Pt-Bi as a potential

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bimetallic alloy for the hydrogen generation reaction through a computational highthroughput screening. This alloy was tested and effectively revealed promising electrocatalytic abilities. It would be interesting to do a crossed comparison of the electrocatalytic abilities of e.g. either Pt-Bi or Pt-Ru in different oxidation reactions. In fact, such study has already been suggested by Fujiwara et al. [36]. The electrochemical oxidation of different fuels in acidic media was examined on eight electrodes: Pt, Ru, Rh, Ir, Pd, Au, PtRu and glassy carbon. Lots of fuels were regarded, among which methanol, ethanol, propan2-ol and formic acid (also ethylene glycol, formaldehyde, glucose, etc). With respect to PtRu, the oxidation current of fuel compounds obtained in cyclic voltammetry showed better performances with formaldehyde, formic acid, ethanol, methanol and propan-2-ol. However the maximum power densities of DLFCs were found following the ranking propan-2-ol, methanol, formic acid, formaldehyde and ethanol. In this study, the Fujiwara et al.’s objective [36] was to propose new liquid fuels for DLFCs like e.g. glucose and potential anode electrocatalysts for these new liquid fuels. This could be completed by attempting to explain these classifications in terms of mechanisms and electronic and geometric effects on the basis of the physical, chemical properties of the chemicals. That might be the purpose of another article. From the observations about both Pt-Pd and Pt-Au, Demirci [14] suggested that other bimetallic electrocatalysts might be found if the following assumptions were considered: Pt antisegregates and the Pt d-band center shifts up. One may add that Pt should not mix in the surface layer [23]. A single choice emerged, Pt-Ag. Yu and Pickup [82] reported Demirci’s [14] DFT-based prediction but well remarked that there has not yet been any experimental study to investigate such hypothesis. On the other hand, unfortunately, Demirci [14] considered only Pt. It may therefore be interesting to do the same work for Pd as the host metal of potential bimetallic alloys. The following assumption is considered: the Pd d-band centre should shift up. If the segregation is not taken into account, alloying Pd with Fe, Co, Ni, Cu, Ru, Rh, Ag, Ir, Pt and Au are some possibilities. If the segregation of Pd is regarded, both Ag and Au are eliminated. No information with respect to CO-induced segregation is available for the Pt- and Pd-based bimetallic alloys [13], which were experimentally investigated and are listed in the previous paragraphs. One might expect weaker CO-bonding over the metals of the alloy. Of course, that has to be confirmed. With respect to trimetallic alloys, combinations of Pt, Pd, Ag and Au, with Pt as the permanent feature, were suggested elsewhere [14]. By using the methodology described by Demirci [14], Pd-based trimetallic alloys could be suggested. Accordingly, combinations of Pd, Fe, Co, Ni, Cu, Ru, Rh and Ir may be proposed. Given what happened with both DMFC and DEFC, investigations about DFAFC will certainly focus on and favor trimetallic combinations.

3.5 Direct Borohydride Fuel Cell Sodium borohydride (sodium tetrahydroborate) NaBH4, as an alkaline aqueous solution, is one of the last chemicals that have been considered as a potential fuel for the DLFC [103]. Up to recent years, NaBH4 has especially been considered as a reducing agent much used in organic chemistry and as bleaching agent in the manufacture of paper. NaBH4 as an energy/hydrogen carrier is in fact not a new compound since it has been discovered in the

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1940s [104-106]. It is an energy carrier when it directly powers a PEMFC-type fuel cell, socalled direct borohydride fuel cell (DBFC) [103, 107]; it is a hydrogen carrier when it stores atomic hydrogen (and generates hydrogen to feed a PEMFC) [108]. In other words, NaBH4 is able to feed directly or indirectly a PEMFC (Figure 9).

Aqueous alkaline solution of NaBH4 Oxidation

Hydrolysis Direct fuel

Indirect

H

Fuel cells

Direct borohydride fuel

H2-powered PEMFC

Figure 9. Borohydride oxidation (blue area) versus borohydride hydrolysis (pink area) for the PEMtype fuel cells.

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In the 1940s, sodium borohydride attracted attention because of its ability to generate molecular hydrogen, the generation reaction being spontaneous and besides being able to be accelerated thanks to either acids or metal-based catalysts like e.g. cobalt or ruthenium [105, 109, 110]. In the meantime, the NaBH4 oxidation began drawing some researchers’ attention [111]. In 1953, Pecsok [112] proposed for the polarographic oxidation of NaBH4 the net reaction (9) and calculated the standard electromotive force of this half reaction as being 1.23 V versus the standard hydrogen electrode: BH4- + 8 OH- Æ BO2- + 6 H2O + 8 e-

(9)

In 1962, Indig and Snider [113] proposed the concept of the DBFC because the possibility of using NaBH4 as an anodic fuel appeared attractive due to its extremely low equivalent weight of 4.73 g. However it was concluded that the possibility of achieving a reversible NaBH4 electrode with an 8-electron oxidation appeared unlikely and that one might have to settle for a reaction involving 4-electron change (10) and a lower potential: BH4- + 4 OH- Æ BO2- + 2 H2 + 2 H2O + 4 e-

(10)

From the 1960s to the late 1990s, NaBH4 was not considered in the energy area anymore. However, since 1999, the year of Millenium Cell’s articles publishing, NaBH4 attracted attention one more time and today increasingly interests as energy/hydrogen carrier. DBFC is based on the borohydride oxidation. The oxidation reaction takes place under an alkaline medium because the borohydride ions are not chemically stable in acidic media. Basically, these ions self-hydrolyze generating molecular hydrogen according to the theoretical reaction (11) that follows:

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NaBH4 + 2 H2O Æ NaBO2 + 4 H2 (11) In order to hinder this undesired reaction, the NaBH4 solution is stabilized with sodium hydroxide. However, in presence of specific metals (e.g. Ni, Ru, Pt…), this reaction is catalyzed and in the DBFC, the borohydride oxidation and the borohydride hydrolysis compete. This is one of the issues hindering the development of the DBFC [103, 114]. The occurrence of the borohydride hydrolysis is detrimental to the DBFC performances since the oxidation does not generate eight electrons per borohydride anymore. For instance, with Ni or Pt electrocatalysts, the oxidation is a four-electron process. The reason of this competition is simple: both reaction media are identical (Figure 10) and for both reactions the electrocatalysts are transition or noble metals [115, 116]. Today the main challenge is to find highly active and selective metal-based electrocatalysts for the DBFC. Some reviews discussed the anode electrocatalysts of the DBFC [103, 108, 114]. For more details about fundamentals, syntheses and performances the reader can refer to the references that are discussed in the present paragraph. Ponce de Leon et al. [114] especially reviewed e.g. the chemical reactions, the fuel cell configurations, the electrode materials for the oxidation of borohydride and the influence of electrolyte parameters. Wee [108] reviewed the two types of fuel cell systems using the NaBH4 aqueous solution as fuel: the hydrogen/air PEMFC (sodium borohydride stores and generates hydrogen) and the DBFC; the performances and the costs of both technologies were particularly discussed. Demirci [103] reviewed the main issues met by the membrane-electrodes-assembly of the DBFC, focusing especially the anode electrocatalyst (metal nature, syntheses, performances, mechanistic aspects, etc.).

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Aqueous alkaline solution of NaBH4 NaBH4  NaOH (pH ~ 13‐14)  H2O  Metal catalyst 

Oxidation medium

Hydrolysis medium

NaBH4  NaOH (pH ~ 13‐14)  H2O  Metal catalyst 

Competition oxidation vs. hydrolysis Figure 10.NaBH4: oxidation and hydrolysis media.

With respect to the metal-based anode electrocatalysts, Demirci [103] listed the metals that were studied: that is, Ag [117], Au [117], Cu [118], Ni [118], Os [119], Pd [118], Pt [118, 120], and bimetallic combinations of Ag [121], Au [122, 123], Pt [122, 123], Pd [123], Sn [119], Os [119], Mo [119], V [119], Ni [122], Ir [122]. Besides it was reported AB2- and AB5-type hydrogen storage alloys as anode electrocatalysts for the DBFCs [124-127]. Today, the main challenge is to find highly active and selective electrocatalysts, that is, electrocatalysts oxidizing borohydride while involving the maximum of eight electrons.

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Investigations are in progress. Among the various metals tested, Au is the most promising. With Au, the borohydride oxidation is ideally a seven- or eight-electron process. Similar performances were reached with Ag that involved 7.6 electrons per oxidized borohydride [117]. For the metals assessed, Demirci [103] reviewed the number of electrons generated per oxidized NaBH4; typically, with Ag and Au, the oxidation involves 7-8 electrons; with Pd 6 electrons; with Pt and Ni 2 to 4 electrons. The d-band center values of these metals are as follows: Au (-3.56 eV), Ag (-4.30 eV), Pd (-1.83 eV), Ni (-1.29 eV) and Pt (-2.25 eV). It was besides reported that, what was expectable, the hydrogen binding energies are weaker for Au and Ag than for Pd, Pt and Ir [25]. In other words, the nobler the metal is, the higher the number of electrons involved in the borohydride oxidation; the less the metal hydrolyses borohydride. While suggesting potential solutions to the issues facing the DBFC, Demirci [103] emphasized that for reaching the ideal eight-electron oxidation, the solution may be the Au-based bi- or tri-metallic electrocatalysts. Besides Au and Ag, no transition metal permits to get the ideal eight-electron oxidation [103]. However, with the noble Au and Ag the kinetics is slow [128]. To enlarge the number of potential electrocatalysts, the bimetallic alloys are conceivable as it is the case in both the DMFC and the DEFC. Gyenge and co-workers [121-123] studied Ag-, Au- and Pt-based bimetallic alloys: that is, Ag–Pt, Ag–Au, Ag–Ir and Ag–Pd [121], Au–Pt and Au–Pd [123], and Pt–Ir and Pt–Ni (and Pt–Au) [122]. Note that Os-alloys (Os–Sn, Os–Mo and Os–V) were also tested without convincing results [119] but the authors yet indicated some activity with significant hydrolysis [128]. Alloying Au or Ag with another metal aims to find electrocatalysts with improved reactivity while permitting an eight-electron reaction. Gyenge and co-workers [128] reported that qualitative tools guided them in the design of their bimetallic electrocatalysts; in the case of Ag-alloys, they used Ag to mitigate the hydrolysis process together with a second metal with a higher activity for borohydride oxidation. Among the bimetallic combinations, Pt–Au was as performing as Au in terms of number of electrons involved in the oxidation since both generated 8 electrons [122]. In such bimetallic materials, while Au enabled the eight-electron oxidation, Pt was expected to dope the electrocatalyst reactivity. For Pt-Au alloys, Au is expected to segregates (Pt antisegregates) and its d-band center should shift up to -3.10 eV from -3.46 eV (no variation for the d-band center of Pt); moreover H-induced segregation is expected [25]. Feng et al. [129] showed that both Ag and Ag-Ni catalyzed the oxidation of borohydride, delivering a high capacity of more than sevenelectron oxidation for a borohydride ion; Ag-Ni-catalyzed borohydride fuel cells exhibited a higher discharge voltage and capacity, possibly due to a combined action of the electrocatalytic activity of Ni component for borohydride oxidation and the depression of borohydride hydrolysis by Ag atoms. In Ag-Ni, Ag strongly segregates and its d-band center shifts up to -3.60 eV (note that this value is that of pure Au). Gyenge and co-workers [128] studied Ag-Au and reported a reaction involving 12 electrons but remarked that such a value that is superior to the maximum of eight electrons required further confirmation. Besides AgAu, other combinations, like Ag-Pt, Ag-Ir and Ag-Pd, were assessed. Such alloying led to a reduction in the number of electrons involved in the borohydride oxidation. According to the theoretical databases, Ag when alloyed to Pt, It or Pd segregates and its d-band center shifts up. This does not agree with the previous observations and therefore it has become difficult to suggest any trend in the reactivity of the bimetallic alloys. While Au-Pt oxidizes borohydride with the involvement of 8 electrons, Ag-Pt would involve only 3.4 electrons. Today there is only one result for each alloy and more experimental results are required to draw trends in the

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alloys reactivity. Furthermore, there is actually no clear evidence about the contribution of the oxidation of hydrogen that stems from the unwanted NaBH4 hydrolysis. The oxidation of this in-situ generated hydrogen should generate 2 electrons per hydrogen; knowing that 4 hydrogen molecules may be formed from one NaBH4, a maximum of 8 electrons can be expected. However one thing is certain: the noble character of the electrocatalyst is crucial (note that no metal when alloyed to either Ag or Au can lead to a down-shift of the d-band center of Ag or Au).

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4. CONCLUSION Computational methods are revolutionizing metal-based catalysis and therefore the searching for potential metallic electrocatalysts. Trial and error has been the traditional method of finding the best catalyst for a reaction, but a computational approach can reduce the lab work required. This comment was Mavrikakis’ [11] while he was discussing Nørskov and co-workers’ DFT calculations-based work [29]. The DFT calculations have provided useful databases for predicting the reactivity of metal-based electrocatalysts: the value of the metal d-band center (and its shift when two metal alloys), the segregation occurrence (for bimetallic alloys) and the mixing of two metals in the surface layer. These literature-available databases have then been used as fundamental tools for understanding the reactivity of monoand bi-metallic materials used as electrocatalysts in various DLFCs. In that way, the methanol-, ethanol-, propanol-, formic acid- and borohydride-powered DLFCs have successively been considered throughout the discussion. We have attempted to interpret the catalytic behaviors of the mono- and bi-metallic anode electrocatalysts. For the oxidation of the alcohols, the best current electrocatalysts are bimetallic alloys: e.g., Pt-Ru and Pt-Sn for the DMFC and the DEFC, respectively. The addition of the second metal (Ru or Sn) improves the reactivity and durability of Pt. It is commonly admitted that because of electronic effects in the bimetallic alloys the adsorption of CO-like species (poisons) over Pt is weakened while the second metal acts as a source of OH species necessary to oxidize the CO-like species adsorbed over Pt. It is the principle of the so-called bifunctional mechanism. Such a mechanism is besides confirmed by the DFT calculationsbased databases. For example, it is predicted that in the methanol oxidation, the Pt d-band center shifts up (i.e., weakening of the adsorbate adsorption) when it is alloyed with Ru and that the surface Pt atoms dilute the surface Ru sites (i.e., small Ru sites providing OH species). Similar interpretations are given for Pt-Sn as electrocatalyst for the ethanol oxidation. With respect to the propanol oxidation, very few papers are available when compared to both methanol and ethanol. Today it seems likely that the best bimetallic materials are Pt-Ru and Pt-Sn, knowing that the potential electrocatalyst must be able to cleave one or two C–C bonds in addition to the C–O bond. Regarding the formic acid oxidation, Pd-based bi- and tri-metallic electrocatalysts have shown the best performances. As it is the case for the bimetallic alloys used for the alcohols oxidation, the bifunctional mechanism explains the improved reactivity of the Pd-based alloys in relation to the pure Pd electrocatalyst. Furthermore, the DFT calculations-based databases have confirmed the electronic and geometric effects explaining the improved reactivity of the alloys. With respect to trimetallic alloys, by using a methodology described elsewhere [14],

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Pd-based trimetallic alloys have been suggested: that is, combinations of Pd, Fe, Co, Ni, Cu, Ru, Rh and Ir, with Pd as the permanent feature. With respect to the borohydride oxidation, the reaction mechanisms are completely different from those of the organic fuels. It has been shown that the borohydride eightelectron oxidation, which is the expected reaction, competes with the borohydride hydrolysis, which is the undesired reaction. The former reaction is favored over Ag and Au electrocatalysts that have very low d-band centers while the latter is observed with highly reactive metals like Ru and Pt that have higher lying d-band centers. In other words, the nobler the metal is, the higher the number of electrons involved in the borohydride oxidation; the less the metal hydrolyses borohydride. Very few bimetallic alloys have been envisaged and therefore no interpretation can be suggested. Globally the present discussion about the metallic electrocatalysts of various DLFCs has shown that the reactivity trends from experimental findings are consistent with the DFT-based theoretical predictions. Thus, the DFT calculations-based approach including the d-band center, the segregation and the surface mixing is thought to be useful to interpret the metal electrocatalysts reactivity and then to search for new electrocatalysts (e.g., trimetallic combinations). According to Mavrikakis [11], the computational approach has great potential to speed up the whole catalysts discovering procedure and leads to better catalysts. And this is effectively the final objective of the DFT even if here it has been very useful for interpreting the reactivity of already discovered efficient, good oxidation electrocatalysts. The computational approach has shown its usefulness downstream from the experimental evidences, but its real power is the possibility to use it as a screening tool upstream from the experimental work.

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[115] Amendola, S. C.; Sharp-Goldman, S. L.; Janjua, M. S.; Kelly, M. T.; Petillo, P. J.; Binder, M. J. Power Sources 2000, 85, 186-189. [116] Demirci, U. B. Int. J. Hydrogen Energy 2008, 33, 2123-2124. [117] Chatenet, M.; Micoud, F.; Roche, I.; Chainet, E. Electrochim. Acta 2006, 51, 54595467. [118] Liu, B. H.; Li, Z. P.; Suda, S. Electrochim. Acta 2004, 49, 3097-3105. [119] Atwan, M. H.; Northwood, D. O.; Gyenge, E. L. Int. J. Hydrogen Energy 2005, 30, 1323-1331. [120] Kim, J. H.; Kim, H. S.; Kang,, Y. M.; Song, M. S.; Rajendran, S.; Han, S. C.; Jung, D. H.; Lee, J. Y. J. Electrochem. Soc. 2004, 151, A1039-A1043. [121] Atwan, M. H.; Northwood, D. O.; Gyenge, E. L. Int. J. Hydrogen Energy 2007, 32, 3116-3125. [122] Gyenge, E.; Atwan, M. H.; Northwood, D. O. J. Electrochem. Soc. 2006, 153, A150A158. [123] Atwan, M. H.; MacDonald, C. L. B.; Northwood, D. O.; Gyenge, E. L .J. Power Sources 2006, 158, 36-44. [124] Wang, L.; Ma, C.; Mao, X. J. Alloys Compd. 2005, 397, 313-316. [125] Lee, S. M.; Kim, J. H.; Lee, H. H.; Lee, J. Y. J. Electrochem. Soc. 2002, 149, A603A606. [126] Li, Z. P.; Liu, B. H.; Arai, K.; Suda, S. J. Electrochem. Soc. 2003, 150, A868-A872. [127] Choudhury, N. A.; Raman, R. K.; Sampath, S.; Shukla, A. K. J. Power Sources 2005, 143, 1-8. [128] Atwan, M. H.; Gyenge, E. L.; Northwood, D. O. Int. J. Hydrogen Energy 2007, 2008, 33, 2125-2126. [129] Feng, R. X.; Dong, H.; Cao, Y. L.; Ai, X. P.; Yang, H. X. Int. J. Hydrogen Energy, 2007, 32, 4544-4549.

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Chapter 8

THE ELECTRO-CATALYTIC AND MASS TRANSPORT COMPONENTS OF THE ELECTRODE POTENTIAL LOSS IN A PAFC CATHODE DETERMINED USING A NUMERICAL MODEL G. Psofogiannakis1,4,***, Y. Bourgault2, B.E. Conway3**, and M. Ternan5*

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The University of Ottawa, 1Department of Chemical Engineering, 2Department of Mathematics and Statistics, 3Department of Chemistry, and 4Centre for Catalysis Research and Innovation, Ottawa, Ontario, K1N 6N5, Canada 5 EnPross Inc., 147 Banning Road, Ottawa, Ontario, K2L 1C5, Canada

ABSTRACT The electro-catalytic and mass transport components of the electrode potential loss in the cathode of a phosphoric acid fuel cell (PAFC) were investigated through the development a numerical model. A one-dimensional model was used to describe the porous electrode that was composed of agglomerates containing platinum catalyst / carbon / polytetrafluoroethylene components. The porous structure of the spherical agglomerates was partially filled with phosphoric acid electrolyte while a liquid film of electrolyte surrounded the agglomerates. The system of equations that described the system was solved using the shooting method. The predictions made by the model were in good agreement with two different sets of experimental polarization data. Increasing the loading of the platinum electro-catalyst from 0.01 to 1 mg/cm2 increased the current density by one and one-half orders of magnitude. Increasing the thickness of the electrocatalyst layer from 50 μm to 250 μm increased the current density by a factor of three. Decreasing the radii of the electro-catalyst containing agglomerates from 5 μm to 0.1 μm also increased the current density by a factor of three. The oxygen reactant concentration profiles, from the agglomerate exterior to its center, were determined. The oxygen *** Current Address: 17 Themistokleous Str., Iraklio, Crete, Greece 71307 ** Deceased * Corresponding author: ternan @ sympatico.ca

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G. Psofogiannakis, Y. Bourgault,B.E. Conway et al concentration at the center was essentially zero in agglomerates having a diameter of 2 μm, whereas it was 73% of the exterior concentration in agglomerates having a diameter of 0.5 μm. In addition, the model identified the fraction of the electrical potential loss caused by each of the following phenomena; electro-catalytic surface processes (cathode overpotential), ionic conduction, diffusion of oxygen gas, and diffusion of oxygen dissolved in the liquid electrolyte. Of these, the electro-catalytic surface processes contributed the largest single electrical potential loss.

INTRODUCTION This work is part of a project that is directed toward the understanding and improvement of phosphoric acid fuel cells (PAFCs) that use hydrocarbons, such as propane, as the fuel directly at the anode [1]. PAFCs that operate using hydrogen as the fuel are a mature technology [2]. The cathode of a PAFC should be similar regardless of whether the fuel at the anode is hydrogen or a hydrocarbon. The performance of the cathode is the subject of the present investigation. Although optimization studies of PAFCs have been performed in the past [3], some of the cathode electro-catalyst properties have not been treated in the detail presented here. This study was specifically intended to investigate cathode electro-catalyst properties, such as platinum loading, thickness of the cathode electro-catalytic layer, and the size of the carbon-platinum agglomerates. In addition the magnitude of the energy losses associated with each of the phenomena were determined. This work describes a mathematical model and compares its predictions to experimental results obtained using a PAFC cathode. The electrochemical oxygen reduction reaction (ORR) at the cathode,

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O2 + 4 H+ + 2 e – = 2 H2O

[1]

takes place within a complex reaction system. Fuel cell cathodes often consist of a planar gas diffusion layer (GDL) that has one of its sides adjacent to the gas channel (GC) providing the O2 reactant. The other side of the GDL is in contact with a catalyst layer (CL). The CL is between the GDL and the electrolyte layer (ELL). The most common cathodes are prepared by applying a very thin layer of catalyst paste onto a porous carbon sheet. The CL contains three components, particles of polytetrafluoroethylene (PTFE), particles of porous carbon, and platinum or platinum alloy catalyst crystallites dispersed within the pores of the carbon particles. PTFE is used to bond the carbon/catalyst particles together and to make the CL more hydrophobic so that flooding by the electrolyte in the CL is restricted and better reactant gas distribution within the GDL is possible. The oxygen reactant gas enters the CL from the GDL and the proton reactant ions that are required for the reaction shown in Equation 1 enter from the ELL. Several mathematical models of fuel cells have been described. Homogeneous models [4,5,6] in which the chemical composition is considered uniform throughout the electrode require the least detail. Non-homogeneous models of the CL geometry have progressively become more comprehensive. The early models [7] that considered three-phase solid / liquid / gas contact were; the simple pore model [8,9,10], porous electrode theory [11], the thin film model [12], and models for electrodes with bimodal pore sizes [13,14,15]. The thin-film

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241

model and the simple pore model were compared by Austin et al. [16]. These models employ simple geometries and include an elementary treatment of the mass-transport phenomena in the electrode. The more complex flooded agglomerate model of the CL was first proposed by Giner and Hunter [17] to describe gas diffusion electrodes. Several variations, extensions and discussions of the flooded agglomerate models appear in the literature ( Cutlip [18], Maggio [19], Kunz and Gruver [20], Iczkowski and Cutlip [21], Ross [22], Perry et al. [23], Bjornbom [24], Yang et al. [25,26], Maoka [27], Yang [28], Choudhury et al. [29] ). In all agglomerate models, including the one presented in this work, the description of the phenomena in the electrode catalyst layer is similar. The catalyst layer is presented as a collection of homogeneously distributed agglomerates of Pt/C catalyst particles having spherical, cylindrical or flat slab geometry. The liquid electrolyte is concentrated within the smaller pores of the agglomerates and some of the models consider it to form a thin liquid film that surrounds the agglomerates. The PTFE particles are hydrophobic, so that the space between the agglomerates consists of gas-filled pores that distribute the reactant gas. In all these models the reactant gas (O2 for this discussion) dissolves in the electrolyte at the exterior surface of the agglomerate or the liquid film and diffuses towards the catalyst sites within the electrode. Thus the model requires the mathematical description of gas diffusion phenomena in both gas-filled and liquid-filled pores, the reaction on the catalyst surface and the ionic transport. The electrical potential gradient required for ionic conductivity in the electrolyte in the catalyst layer and the oxygen concentration distribution within the catalyst layer are responsible for the development of both electrode potential and current density distributions throughout the catalyst layer. The present model and the results obtained using it are more comprehensive than any of the previous flooded-agglomerate models for phosphoric acid cathodes. First, there are three unique attributes of this work compared to previous models. The model includes (a) the combined effects of temperature, pressure, and H3PO4 concentration on the oxygen diffusivity and on the oxygen solubility, (b) the effect of temperature on the reaction rate, (c) the effect of catalyst composition (%Pt in Pt/C) on the effective platinum surface area, (d) the effect of temperature on the water vapor pressure, and (e) the effect of temperature on the equilibrium potential. Other effects that have been included in previous models are also included in this model, such as the dependence of the catalyst utilization on the electrolyte fill-up ratio in the catalyst layer [19]. The additional effects render temperature, pressure, acid concentration and catalyst composition as independent variables in the model. Previous agglomerate models did not examine the effects of these variables and thus required the specification of a consistent set of variables (for example for diffusivity and reaction rate, which both depend on temperature). In this model, all other model parameters can be estimated from these as well as from the micro-structural characteristics of the electrode. Other unique features include the derivation of the reaction rate equation with consideration of spherical shell liquid film geometry and different dissolved gas diffusivities for the agglomerate and the liquid film. Second, the present model estimates the relevant morphological parameters of the catalyst layer from a set of natural, measurable, electrode characteristics. The approach is unique in that many morphological parameters, particularly, the volume fraction of agglomerates, the agglomerate porosity, the gas volume fraction, the catalyst utilization, the catalyst available surface area and the liquid film thickness can be estimated from the following electrode characteristics: catalyst layer thickness, platinum loading, catalyst

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composition, H3PO4 fill-up ratio in the catalyst layer, and agglomerate size. Another unique feature is that this model also allows for the possibility of dry agglomerates if the electrolyte content of the electrode (the electrolyte fill-up ratio) is insufficient to flood all the pores that are interior to agglomerates. Previously Yang [28] included a morphological model for the electrode. However, it required the specification in advance of the electrocatalyst utilization, a non-measurable parameter. Maggio [19] performed a sensitivity analysis on his model to investigate the effects of several model parameters, some of which are related to the morphology of the electrode (catalyst utilization, percentage of acid (PAO) in the catalyst layer, Pt loading in the catalyst layer, catalyst surface area, fraction of Pt in the catalyst layer). The different approach in the present model is that the interdependence between all model parameters is considered. For example, catalyst utilization is a function of both the PAO and the catalyst surface area, which in turn is a function of the catalyst composition (fraction of Pt in the Pt /C particles). Third, many of the results obtained from this study are unique. As discussed above, this model is capable of simulating electrodes of different characteristics and at different operating conditions. The results include the effects of variables not included in other models of phosphoric acid cathodes such as temperature, acid concentration, pressure, catalyst composition in addition to those variables included in other models, such as catalyst layer thickness and platinum loading. Furthermore, we have calculated the relative magnitude of electrical potential losses caused by different phenomena in the cell. Specifically, the electrical potential losses caused by reaction, reactant diffusion in the liquid and gas phases and ohmic conduction were determined as a function of both the current density and the partial pressure of oxygen at a specific position in the cathode gas channel. As the current density increases, the electrical potential losses caused by the diffusion processes increase. Also, at a position close to the entrance of the gas channel where the oxygen partial pressure (related to reactant utilization) is large, the effect of reactant diffusion limitations would be expected to be smaller than at a position further downstream where reactant depletion causes larger electrical potential losses (concentration polarization). Previously, Yang et al. [26] discussed the effects of current density and reactant utilization (local). They presented the distribution of potential losses, resulting from their model, at a particular oxygen utilization for five different current densities and at two different oxygen utilizations for one current density. In contrast, this work describes these effects for both the entire current density range and the entire oxygen utilization range. Furthermore, contour diagrams were used to present the combined effect of both variables, as well as the contributions of the processes that cause electrical losses at all operating conditions.

MODEL DESCRIPTION The principal features of the isothermal model used in this development are reported here. Details such as the computer code have been documented elsewhere [30]. The model is based on the following eleven assumptions: (1) The composition of gas at any point in the gas channel is a homogeneous mixture of reactant (oxygen or air) and product (water) that corresponds to a specific oxygen conversion

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243

within the range 0% to 100%. For this one-dimensional model, the oxygen partial pressure in the gas channel is specified. (2) The electrode can be divided into two separate layers: the gas diffusion layer and the catalyst layer. The gas diffusion layer contains no liquid electrolyte. (3) The catalyst layer consists of spherical agglomerates that contain all of the carbon support and dispersed platinum catalyst. All particles of PTFE are exterior to the agglomerates [31]. The pore volume within agglomerates can only be either completely filled with electrolyte or completely filled with gas. The existence of dry agglomerates is accounted for because there may not be enough electrolyte in the catalyst layer to wet all the agglomerate interior pores, in which case some of the platinum is not utilized. Agglomerates that are completely filled with electrolyte are utilized because they can contribute to the transferred current. Thus for the present model, the under-utilization of catalyst in the electrode is the combined effect of both non-wetted catalyst (dry agglomerates) and the agglomeration of platinum crystallites, which is considered through the dependence of platinum surface area [m2/(g Pt)] on the catalyst composition (weight fraction of platinum in Pt+C). (4) The space between agglomerates consists of gas-filled pores and particles of PTFE. The largest pores in the catalyst layer are exterior to agglomerates and the smaller pores are interior to agglomerates. This suggests that capillary forces will act to fill the agglomerate pores first. The pores within agglomerates will be referred to as interior pores. There also exist true micropores that are within the individual carbon particles of large surface area carbon supports. Watanabe et al. [32] have noted that these micropores are too small to be suitable either as a diffusion path or as a surface for accessible reaction sites. Therefore, the micro-structural description of the electrode in this model does not include these micropores. (5) The percentage of utilized agglomerates is linearly related to the electrolyte fill-up ratio FA(fraction of pore volume in the electrode that is occupied by liquid). At a particular electrolyte fill-up ratio FA denoted as FA,MAX, all of the agglomerates are utilized. Another term for FA is the percentage of acid occupation (PAO). The percentage of interior pores filled with electrolyte is equal to the agglomerates’ utilization. Justification for this assumption is given by Giordano et a1. [31], in which it is also shown that FA,MAX is approximately equal to 0.7. (6) All utilized agglomerates are covered by thin electrolyte films. Agglomerates that don't contain electrolyte are not covered by liquid films. The effective electrolyte film thickness can be estimated from knowledge of the FA,MAX and the agglomerates radii rAG. It is assumed that the films can grow thicker if more acid is contained in the catalyst layer than that which fills all the micropores. Thus, this concept allows for performance deterioration due to flooding of the electrode. (7) All agglomerates are homogeneously distributed in the catalyst layer, contain the same amount of carbon and platinum and have the same radii. (8) The solid phase does not present an appreciable resistance to the flow of electrons through it. (9) The water produced by the electrochemical reaction is transferred out of the electrode pores so that the phosphoric acid concentration does not change during the transport of current. The water produced by the reaction is assumed to flow towards the gas channel (and not towards the electrolyte layer) at the same rate that it is produced. This is a steady-state model. In reality, the water evaporation rate maybe smaller than the water production rate or

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the produced water may diffuse towards the electrolyte layer, in which case the acid concentration will decrease with time. Such effects are not considered in this model. (10) The model considers a double distribution of oxygen reactant in the electrode: gasphase oxygen concentration is distributed axially across the thickness of the electrode. Dissolved oxygen concentration is distributed radially within each individual agglomerate. The agglomerate diameter is negligible compared to the catalyst layer thickness, so that each agglomerate is assumed to correspond to a particular gas-phase oxygen concentration. Phase equilibrium is assumed to relate the gas-phase oxygen concentration to the dissolved oxygen concentration at the exterior surface of the liquid film surrounding each of the utilized agglomerates. Similarly, the overpotential varies across the thickness of the CL since it is related to the electrical potential in the electrolyte that also varies across the CL and is considered to be the driving force for proton transport into the agglomerates. However, for each individual agglomerate, at a particular position in the CL, the overpotential is assumed to be constant (it does not vary radially inside agglomerates). (11) The flow mechanism in the gas-filled electrode pores across the thickness of the GDL and the CL is taken to be only molecular diffusion. The diffusion fluxes caused by Knudsen diffusion were calculated to be negligible compared to molecular diffusion for the range of pore radii of the gas diffusion layer and catalyst layer [33]. The input parameters that are used in the model are the thickness of the catalyst layer LCL (cm), the thickness of the gas diffusion layer LGDL, (cm), the weight fraction of platinum in the catalyst WF/Pt [g Pt /g (Pt+C)], the platinum loading per unit electrode face surface area MPt (g Pt /cm2), the fraction of pore volume filled with electrolyte in the catalyst layer FA, the fraction of total pore volume that is interior to agglomerates in the catalyst layer (FP)IN, the weight fraction of polytetrafluoroethylene in the solid in the catalyst layer WF/PTFE (g PTFE I g (PTFE+Pt+C)), the weight percentage (concentration) of the phosphoric acid electrolyte W (%), the operating temperature T (K), the gas pressure P (Pa) and the composition (mole fraction of oxygen) of the reactant gas. These parameters can be specified or experimentally measured for a specific electrode. The flow in the gas-filled pores in the CL and the GDL is described by the StefanMaxwell molecular diffusion equations. According to the Stefan- Maxwell equations [34,35] for the ternary gas system O2, N2, and water vapor (W), the corresponding diffusive fluxes and concentrations are related by: . – dcO2G / dz = [cWG NO2 – cO2G NW] / [cT DEO2,W] + cN2G NO2 / [cTDEO2,N2]

[2]

. – dcN2G / dz = cN2G NO2 / [cT DEO2,N2] + cN2G NW / [cTDEN2,W]

[3]

where z is the distance coordinate along the thickness of the CL or GDL, cT is the total concentration of all the gaseous components in the gas phase, and DEAB is the effective binary diffusivity between components A and B. Since oxygen does not react within the GDL, its flux will be invariant with position and therefore for the GDL: dNO2 / dz = 0

[4]

In the catalyst layer, the oxygen flux and the water vapour flux are related to the oxygen volumetric reaction rate by the stoichiometry of the reaction. For the CL: Polymer Electrolyte Membrane Fuel Cells and Electrocatalysts, Nova Science Publishers, Incorporated, 2009. ProQuest Ebook Central,

The Electro-Catalytic and Mass Transport Components of the Electrode Potential dNO2 / dz = RV = – ½ [ dNW / dz]

245 [5]

where Rv denotes the volumetric reaction rate in the electrode's catalyst layer (moles O2 consumed / cm3 s). The binary diffusion coefficients, for different temperatures and gas pressures, were estimated using the following empirical equation [36] DAB = [ 0.00143 T1.75 ] / { P(MAB)1/2 [(ΣV)A1/3 + (ΣV)B1/3]2 }

[6]

where DAB is the binary diffusion coefficient of species A in B [cm2/s], P is the gas pressure [bar], T the temperature [K], MAB = 2(MA –1 + MB-1)-1, M is the molecular weight, and (ΣV) is the molecular diffusion volume (18.5 for N2,16.3 for O2, and 13.1 for H2O)36. The reaction rate equation was developed next, as an expression for the volumetric reaction rate in the CL, Rv [moles/(s cm3CL], as a function of the gas-phase oxygen concentration. Because the reactant gas concentration and the overpotential are variable across the electrode thickness, the reaction rate will also be variable with coordinate position, giving rise to a current distribution in the electrode. The starting point is the Butler-Volmer equation, which gives the current density per unit area of platinum, j [A/(cm2 of platinum metal)] as:

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j = j0 [ cO2D / c**] (exp[ – αC zE F η / RT] – exp [ αA zE F η / RT] )

[7]

where j0 is the exchange current density for oxygen reduction, cO2D is the concentration of oxygen dissolved in the electrolyte [mole/cm3] and is varying with coordinate position r across the agglomerate., c** is the solubility of oxygen in the liquid electrolyte at one atmosphere of oxygen pressure [mole/cm3], η is the overpotential [V], αc is the cathodic transfer coefficient for the ORR and αA is the anodic transfer coefficient for the ORR. The current transferred across the metal solution interface per unit volume in the agglomerate is given by jSA where SA [cm2 Pt / cm3 agglomerate] denotes the platinum surface area per unit volume in the agglomerate. Utilizing Faraday's law, the reaction rate per unit volume of agglomerate, RV-AG, is expressed as jSA / zE F, so that: RV-AG = (SA j0 cO2D /zE F) (exp[ – αC zE F η / RT] – exp [ αA zE F η / RT] )

[8]

The treatment of simultaneous diffusion and catalytic reaction in a porous sphere has been described previously in the literature [37]. A schematic diagram of the situation is presented in Fig. 1. The reaction-diffusion equation for an agglomerate can be derived by a mass balance through a differential spherical shell at position r within the agglomerate. The molar flux of dissolved oxygen gas, in the radial direction, within the agglomerate is given by Fick's Law using an effective diffusivity, DE within the porous C-Pt agglomerates. The following equation can be derived [38]: RV-AG / DE = [ 1 / r2 ] d/dr[r2 (dcO2D/dr) ] where RV-AG is given by Equation 8.

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[9]

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Figure 1. Diagram of a porous C-Pt agglomerate showing geometry, nomenclature, and boundary conditions

Equation 9 can be integrated [37] using two boundary conditions: (1) at the centre of the agglomerate, r = 0, the flux of dissolved oxygen is zero (by symmetry) so that d cO2D / dr = 0, and (2) at the interface between the exterior of the C-Pt agglomerate and the liquid film of electrolyte that surrounds it, rAG, the concentration is denoted as cO2D/IN (the interfacial concentration). The solution to the equation [37] is: cO2D = cO2DIN [rAG sinh ( r √ψ )] / [ r sinh ( rAG √ψ ) ]

[10]

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where: ψ = = (SA j0 /zE F c** DE ) (exp [ – αC zE F η / RT] – exp [ αA zE F η / RT] )

[11]

The reaction rate within an agglomerate, RAG (moles s –1 agglomerate –1 ), is equal to the oxygen molar flowrate into the agglomerate across the exterior surface of the agglomerate. RAG = [DE (dcO2D/dr)| r = rAG] (4πrAG2) = 4πDE cO2D/IN rAG [ rAG √ψ coth(rAG √ψ) – 1 ] [12] Another expression for RAG was obtained by considering the diffusion of dissolved oxygen through a spherical shell of electrolyte film surrounding the agglomerate. Since no reaction takes place in the liquid film, the differential mass balance at position r in the liquid film (r > rAG) gives: d/dr [ r2 dcO2D/dr ] = 0

[13]

Equation 13 is subject to the boundary conditions: At r = rAG, cO2D = cO2D/IN and at r = rAG+LF, cO2D = K cO2G. In the last boundary condition, a Henry's law equilibrium relationship

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was used to relate the gas-phase and dissolved oxygen concentrations. The solution of Eq. 13 that satisfies the specified boundary conditions is (r > rAG): dcO2D / dr = [ rAG (rAG + LF) ( K cO2G – cO2D/IN) ] / ( LF r2)

[14]

where K is a Henry's law constant. The value for the reaction rate, RAG, can be found from the flow rate of reactant at the exterior of the liquid electrolyte film: RAG = [ 4 π D rAG (rAG + LF ) ( K cO2G – cO2D/IN) / LF

[15]

where the true diffusivity D was used for the liquid film because it contains only electrolyte and cO2G, is the oxygen concentration in the gas phase that varies through the CL. Equations 12 and 15 can be equated and solved for (cO2)D/IN. The value of (cO2)D/IN can be substituted into Equation 11 to obtain the reaction rate per agglomerate. By multiplying by ρμAG, the density of the agglomerates filled with liquid electrolyte (utilized) (the number of electrolyte filled agglomerates / cm3,) the reaction rate per unit volume of catalyst layer, Rv, can be obtained. RV = ρμAG {4πDE D (rAG + LF) K cO2G rAG [ rAG √ψ coth (rAG √ψ) – 1 ] } / D (rAG + LF) + [16] LF DE [ rAG √ψ coth (rAG √ψ) – 1 ]

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where ψ is defined in Eq. 11. The physical parameters contained in Eq. 16 and 11 were determined by developing correlations from existing experimental data. The following correlation for the variation of the oxygen solubility in phosphoric acid, c*, [(mole/cm3)LIQ / (atm)GAS], with weight percent phosphoric acid concentration, W, was found to fit the experimental data [39] for temperatures greater than 100°C, which is normally the operating temperatures of PAFC's: c* = 10 –7 (6.87 – 0.062 W)

[17]

For operation at temperatures greater than 100°C, the solubility remains practically independent of temperature [12]. Henry's law constant, K is related to the solubility: K = RTc*

[I8]

The diffusivity of dissolved oxygen, D, in the electrolyte was correlated to the phosphoric acid concentration (weight percentage) W% and the temperature T, by combining the StokesEinstein equation relating diffusivity to viscosity with both diffusivity data [40] and viscosity data [41] to obtain: D = 0.048 exp[ – 25770 / RT + 0.034(98 – W) ]

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[19]

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All gas phase diffusivities, DAB, and liquid phase diffusivities, D, were corrected for the porosity and tortuosity of the relevant layer, using the appropriate CL volume fraction to obtain effective diffusivities. For example, the effective binary diffusivity between oxygen and water was determined as follows: DEO2,W = (VG)1.5 DO2,W

[20]

where VG is the volume fraction of the CL composed of gas-filled pores. The model equations pertaining to the micro-structural description of the catalyst layer will be described next. The volume fraction occupied by solid (VS) in the catalyst layer is the sum of the volume fractions of the three solids (platinum, carbon and PTFE) and is given in Eq. 21 as a function of the normally specified electrode characteristics: the platinum loading (MPt), the weight fraction of Pt in the Pt/C catalyst (WF/Pt), the weight fraction of PTFE in the catalyst layer (WF/PTFE) and the thickness of the catalyst layer (LCL). The densities of each of the solids used in the electrode are ρPt = 21.4 g/cm3, ρC = 2 g/cm3 and ρPTFE = 2.2 g/cm3. The agglomerates consist of carbon, platinum and the interior to agglomerates pores. The sum of the volume fractions of these gives the volume fraction of the catalyst layer occupied by agglomerates VFAG in the catalyst layer. VFAG = (VPt + VC) + (FP)IN ( 1 – VS)

[22]

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where (FP)IN denotes the fraction of the total pore volume that is interior to agglomerates. The fraction of agglomerate volume that is occupied by interior pores (agglomerate porosity) is symbolized as VP/AG and given by the ratio of the volume fraction of interior pores in the catalyst layer to the volume fraction of agglomerates. VP/AG = [ (FP)IN (1 – VS) ] / VFAG

[23]

The fraction, u, of agglomerates that are utilized (filled with electrolyte) is related to the fraction of pore volume occupied by liquid, FA (sometimes denoted as PAO, percentage of acid occupation or electrolyte fill-up ratio). Because of assumption 5 described before: u = FA /FA,MAX for FA < FA,MAX

[24]

If FA > FA, MAX, u=l. The liquid volume fraction in the catalyst layer is then given as: VL =-FA (1 – VS )

[25]

The number density of utilized agglomerates (agglomerates per unit volume) is given by the density of agglomerates multiplied by the fraction of utilized agglomerates.

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249 [26]

The volume fraction of the catalyst layer occupied by liquid films (VLF) can then be estimated as it corresponds to the total liquid volume fraction diminished by the volume fraction of liquid that is within agglomerates. VLF = VL – μ VFAG VP/AG

[27]

The thickness of the liquid films can then be estimated. The volume fraction of liquid films in the catalyst layer is equal to the volume-fraction of utilized agglomerates multiplied by the ratio (volume of film/ volume of agglomerates): VLF = [ (u VFAG) { (rAG + LF) – (rAG)3 } ] / (rAG)3

[28]

Rearranging Eq. 28: LF = rAG [ (VF / { u VAG } + 1 ) ]

[29]

Furthermore, if the surface area per unit mass of platinum is denoted as SPt (and is dependent on WF/Pt), then the surface area of platinum per unit volume of agglomerate (SA), to be used in Eq. 11, is given by:

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SA = ( MPt SPt ) / ( LCL VFAG )

[30]

The platinum surface area SPt (cm2 Pt / g Pt) is a function of the catalyst composition WF/Pt [g Pt/ g (Pt +C)]. A correlation of commercial catalyst data [30] was used to describe this important effect (as the Pt content of the catalyst increases, the effective surface area of platinum decreases because of agglomeration of platinum crystallites). SPt = 104 (131 WF/Pt2 – 302.2 WF/Pt + 168.3 )

[31]

The electrochemical parameters in Eq 11 were taken from the literature and the variation of exchange current density with temperature was considered. The exchange current density at 25°C) (j0)298 for the oxygen reduction reaction (ORR) on platinum in phosphoric acid, was reported [42] to be 3.8*10 –13, A/cm2 Pt. The value of j0 at other temperatures was obtained from an Arrhenius-type equation j0 = (j0)298 exp (ΔGΑ / RT)[ 1/298 – 1/T ]

[32]

where ΔGΑ is the activation energy (difference in Gibbs energy between the transition state and the reactant) for the exchange current, having a value [42,43] of 22 kcal/gmol. For potentials lower than approximately 900 mV, the Tafel slope, 2.303 RT/F, has been reported [20] to be approximately 60 mV/decade at 298 K, which corresponds to a value of the cathodic transfer coefficient αc = 1 that was used in this model. The Tafel slope has been found to be independent of electrolyte concentration [43].

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The remaining variable to be specified is the overpotential, η. Its value at any position in the catalyst is given by: η = φS – φL – EEQ

[33]

where φS is the electrical potential in the solid electrode and φL is the electrical potential in the liquid electrolyte in the CL. The electrical potential in the electrolyte is distributed through the thickness of the catalyst layer. The potential distribution in the electrolyte is required for the conduction of protons. The liquid potential is determined by the reaction requirements for protons and therefore its distribution is determined by the reaction rate. The solid potential is assumed constant in the model, because the solid has been taken as infinitely conductive to electrons. Thus the overpotential will also be variable through the catalyst layer. The following expression [40] was used for the variation of the equilibrium potential with temperature, T, for the oxygen reduction reaction (ORR): EEQ = [ – 4.184 ( – 70650 – 8T ln(T) + 92.84 T ) ] / 2 F

[34]

The potential in the liquid electrolyte within the catalyst layer, φL, is related to the net ionic current density, j+ [A/cm2] by Ohm's law, shown in Eq. 35.

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j+ = – kE dφL/dz

[35]

Ross [22] has noted that the use of Ohm's law is a simplification. In the general case, ionic transport can be attributed to the combined effect of ionic conduction and ionic diffusion. Such an analysis requires the use of dilute solution theory, as done by Perry et al. [23], because concentrated solution transport theory would make the model too complex and requires the use of parameters that are difficult to measure. However, in phosphoric acid fuel cells concentrated electrolyte is used. Ohmic conduction, which uses an effective conductivity to describe ionic mass transport as a result of an electrical potential driving force, will produce more realistic results than infinite dilution approximations of the combined diffusion and migration treatment. Furthermore, in the present work the effect of electrolyte concentration on cathode performance is considered. This effect cannot be handled with infinite dilution approximations. The ionic current density, that is used to represent proton mass transport, is related to the volumetric reaction rate in the CL according to: dj+/dz = – 4 F RV

[36]

The following correlation was found to fit the experimental data on the ionic conductivity, k, of greater than 75 wt%, phosphoric acid solutions at various temperatures [44]. k = (8.53*10 –5W2 – 0.0184W + 1.034) [1 + (4.44*10 –6W2 – 1.01*10 –4 W + 0.0114) T – 298] [37]

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The effective conductivity of phosphoric acid in the liquid-filled pores, kE, was related to the above value by means of the following correlation, which corrects for the porosity and tortuosity of the conduction path: KE = (VL) k

[38]

COMPUTATIONAL METHOD A system of differential-algebraic of equations describes the performance of the electrode. The coupling between the variables, whose variation is described by the differential equations in the catalyst layer, can be seen in the following set of equations that illustrate the functional dependencies among them: From Eq. 16, 11, 33, and 5 dNO2/dz = f1 (cO2, φL)

[39]

from Eq. 2 dcO2 /dz = f2 (NO2, cO2, cN2) from Eq. 3 dcN2 /dz = f3 (NO2, cO2, cN2)

[40] [41]

from Eq. 35 dφL / dz = f4 (j+)

[42]

from Eq. 16, 11, 33, and 36 dj+/dz = f5 (cO2, φL)

[43]

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For all of the equations, the coordinate system is specified in the direction from the gas channel toward the ELL. The nomenclature used to specify boundary conditions and the relevant variables in the differential equations are shown in Fig. 2.

Figure 2. Nomenclature used in specifying the boundary conditions and the relevant variables in the differential equations

A complete set of boundary conditions is required to solve the system of Eq 39 – 43 for the gas diffusion layer and catalyst layer. The electrical potential in the electrolyte was specified arbitrarily at z1+, making it a relative quantity: φL (z1+) = 0 [44]

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The concentrations of the gases in the gas channel (GC) were calculated in this onedimensional isothermal model by specifying the total gas pressure and the composition (mole fractions) in the gas channels. Furthermore at the boundaries z0 and z1, the gas concentrations and fluxes are required to be continuous. Therefore: cO2 (z0+) = cO2 (z0–) = cO2/GC

[45]

cN2 (z0+) = cN2 (z0–) = cN2/GC

[46]

cO2 (z1+) = cO2 (z1–)

[47]

cN2 (z1+) = cN2 (z1–)

[48]

NO2 (z1+) = NO2 (z1–)

[49]

Also at the boundary between the CL and GDL, the ionic current density must be zero because liquid-filled pores terminate at this position. j+(z1+) = 0

[50]

Also at the boundary between the CL and the ELL, the oxygen flux in the gas pores is required to be zero because the gas pores terminate at this boundary

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NO2 (z2–) = 0

[51]

A numerical difficulty was encountered in solving the set of differential equations, Eq. 39 – 43, because the boundary conditions can only be specified at the opposite ends of the electrode. The oxygen and nitrogen gas concentrations can only be specified at the boundary, z0+, between the GC and the GDL (Eq. 45 and 46). In contrast the oxygen gas flux must satisfy the zero flux conditions at the boundary, z2–, between the CL and the ELL (Eq. 51). If the differential equations are solved in the forward direction from z0 to z2 using a timestepping ordinary differential equation solver, the O2 and N2 concentrations are specified in the gas channels. Equations 2 and 3 have initial conditions at z = z0, and Eq. 5 has a final condition at z = z2, that the O2 flux must satisfy the zero flux condition at the boundary. This complication prohibited the numerical solution by a one shot-method proceeding in one direction in the electrode. The difficulty cannot be resolved by solving in the backward direction since two final conditions would arise. If the system of equations is solved sequentially for the GDL and CL proceeding from boundary z0 to z2 by arbitrarily choosing a boundary value for NO2 at z0+ then Eq. 51 will not necessarily be satisfied. Also, if the system is solved in the opposite direction by arbitrarily selecting cO2 and cN2 at z2–, then Eq. 45 and 46 will not necessarily be satisfied. Therefore an iterative method was required to satisfy the boundary conditions. The iterative method used to overcome this difficulty is referred to as the 'shooting' method [45]. Application of this method can be described as follows: the missing condition at one end of the system boundary (NO2 at z0+) was used as a control variable in order to satisfy the target condition at the other end of the system boundary (NO2 = 0 at z2–). Initially, the

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control variable NO2(z0+) was set to an arbitrary value and the system of equations was solved sequentially for the gas diffusion and catalyst layers proceeding in the positive z-direction (in which case all equations have specified initial conditions). When the system was solved, the target condition NO2(z2+) = 0 was not necessarily satisfied, i.e. NO2(z2–) ≠ 0. In the next iteration a different value of the control variable was calculated using the Newton-Raphson formula (superscript i refers to the iteration number). NiO2(z0+) = Ni –1O2(z0+) – { Ni –1O2(z2+) / [dNi –1O2(z2–) / dNi –1O2(z0+)] }

[52]

The derivative [dNO2(z2–) / dNO2(z0+)] was evaluated numerically by introducing a small perturbation ΔNO2(z0+) and resolving the system. A new value of NO2(z2–) was obtained and the value of ΔNO2(z2–) i –1 / ΔNO2(z0+) i –1 was an approximation to the derivative [dNO2(z2–)] / [dNO2(z0+)]. It was found by trial and error that a perturbation value of ΔNO2(z0+) = 10 –9 NO2(z0+), produced a rapid convergence rate. The procedure was repeated using the updated value of the NO2(z0-) as the boundary condition for Eq. 4 and the system resolved. This method resulted in satisfying the target condition NO2(z2–) = 0 to within the specified tolerance after very few iterations. The tolerance specified in this simulation was normalized with respect to the inlet flux at z0+. The condition that was to be satisfied for convergence was:

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| NO2(z2–) / NO2(z0+) | < 10 –7

[53]

Within each iteration, the systems of differential-algebraic equations in the GDL and CL were solved using standard differential equation solvers provided in the software package MATLAB. The solver used was an implementation of the Runge—Kutta [46] method. The code was written in MATLAB. For each point on the polarization curve, the solution of the model resulted in a spatial distribution for the variables, cO2, NO2, RV, j+, and ECL = φS – φL, across the gas diffusion and catalyst layers. Fig. 3 is representative of these distributions and was obtained as a result of solving the set of equations for a specific case. The variables have been scaled according to their maximum values in the electrode to facilitate the graphical representation of their spatial distributions.

Figure 3. Distributions of oxygen flux, NO2, gaseous oxygen concentration, cO2, reaction rate, RV, ionic current density, j+, and electrical potential, φL, in the gas diffusion layer and in the catalyst layer.

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In the polarization curve describing the performance of a single electrode, the measured half cell potential (experimental) corresponds to the difference between the solid phase potential, φS, measured at a convenient location in the electrode (usually z = z0 in Fig. 2 and 3), and the electrolyte potential at the electrolyte side of the electrode, φL, (boundary z = z2 in Fig. 2 and 3). In this model it has been assumed that the electronic conduction in the solid is infinitely fast (much faster than ionic conduction). Therefore, the electrical potential in the solid, φS, will be the same throughout the CL, although the electrical potential in the electrolyte, φL, will vary throughout the CL. The electrode potential difference φS – φL, at the CL-ELL boundary (z2) coincides with the measurable electrode potential of a half-cell, shown as the y-axis of the polarization curve. The x-axis of the polarization curve is the current density per unit electrode face area. This quantity can be found by numerically integrating the current produced in the catalyst layer with respect to the coordinate z, from z1 to z2. Thus the simulated polarization curve is the following: Z2

E = [φS – ( φL |z=z2 ) ] versus I = zE F

∫

Z1

RV dz (Definition of the half cell polarization

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curve) In order to simulate the whole polarization curve the following approach was followed: The solid potential, φS, was initially set to a value close to the equilibrium potential EEQ, of the half-cell reaction and the system of equations describing the electrode was solved numerically. Because the electrolyte potential had been specified as a boundary condition, the solid potential had to be specified at the beginning of the simulation. The quantities required to obtain one point on the polarization curve were then computed. Subsequently, the solid potential φS was stepped down to a new value and the procedure was repeated to obtain a new point in the polarization curve. The set values for φS were closely spaced together in order to increase the number of simulated points. In this manner, the whole half-cell polarization curve was simulated. The values for the various electrochemical and physico-chemical properties are described elsewhere [30].

RESULTS AND DISCUSSION The predictions from the mathematical model were compared to experimental data for the half cell electrodes reported by Kunz and Gruver [20], denoted here as Electrode A, and by Maoka [27], denoted here as Electrode B. The experimentally measured characteristics of the two electrodes and their operating conditions are shown in Table 1. These values were used as input to the model in order to simulate polarization curves for operation in oxygen and in air. Two additional characteristics of the cathode were also used as input data; the agglomerate radius and the exchange current density, j0. For Electrode A the value for the agglomerates' radius used in the simulation was specified to be 1 μm and the exchange current density at 25°C (j0)298, was specified to be 5 x 10 – l 3 amp/cm2. The value of 92 kJ/mole, taken from the literature [42,43], was used for the activation energy. Even though these two particular numerical values were used for the

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agglomerates' radius and for the exchange current density, they are both within the range of values that would be expected from physical measurements. The results of the simulations for Electrode A, shown as continuous lines, are compared with experimental data [20] in Fig. 4. The model predictions for both air and oxygen are in excellent agreement with the data.

Figure 4. Model prediction of polarization curves for Electrode A, cathode potential (V) versus current density (mA/cm2), experimental data [20] (circles and squares represent oxygen and air respectively) and model predictions

TABLE 1. CATHODE PROPERTIES

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Electrode

A (ref 20)

B (ref 27)

Temperature [°C]

160

190

Electrolyte [% H3PO4]

96

105

Pt loading [ mg Pt/(cm2face area) ]

0.25

0.63

Catalyst layer composition Carbon type Pt/(Pt+C) [ wt% ] PTFE/(PTFE+Pt+C) [wt%]

Vulcan XC-72 15

5 50

Catalyst layer thickness [μm] Gas diffusion layer thickness [μm]

Vulcan XC-72

50 75

200

200 600

For Electrode B the value for the agglomerates' radius used in the simulation was specified to be 4 μm and the input for the exchange current at 25°C, (j0)298, was 1 x 10–13

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amp/cm2 (with an activation energy of 92 kJ/mole). Again, these values are within the range of values that would be expected from physical measurements. The results of the simulations for Electrode B, shown as continuous lines, are compared with experimental data [27] in Fig. 5. The model predictions are in excellent agreement with the data for air. For oxygen there is a discrepancy between the prediction and experimental data at large current densities, although agreement is excellent at small and medium current densities.

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Figure 5. Model prediction of polarization curves for Electrode B, cathode potential (V) versus current density (mA/cm2), experimental data [27] (circles and squares represent oxygen and air respectively) and model predictions.

The discrepancy between the model-predicted and experimental data for oxygen feed at high current densities has been noted before in literature describing the flooded-agglomerate model. According to Maja et al [6], this discrepancy should be expected because it is attributed to an inherent fault of the representation of the physical microstructure of the electrode in the flooded-agglomerate model. It is interesting to note that a discrepancy similar to that in Fig.5 at high current densities for oxygen feed was predicted [6], based on the flooded-agglomerate model of Giner and Hunter [17]. The authors note that the difference in potential between air and oxygen feeds is overestimated at high currents because of the assumption of uniform catalyst layer physical dimensions. For the Giner and Hunter model, this pertains to the size of agglomerates as well as the thickness of the catalyst layer. Because, in reality, these dimensions are variable in the catalyst layer, the concentration distributions (for dissolved oxygen in the agglomerates and gaseous oxygen in gas-filled pores) are not well described by a model that is based on a constant value for these dimensions. This results in overestimation of the polarization difference between air and oxygen feeds and the limiting current density for the higher oxygen partial pressure feeds. A second possible reason might be that the oxygen solubility for high oxygen content in the gas-phase could actually be lower than predicted in the model using Henry's law equilibrium condition between gaseous and dissolved oxygen (if there is a departure from equilibrium or if the solubility does not vary linearly with partial oxygen pressure, as taken in the present model). Nevertheless, two observations can be made regarding these results: firstly, the agglomerate model is capable of predicting mass-transport limited data for not too concentrated oxygen feeds (air feed). This is

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shown by the predicted polarization curves in this work as well as in Maja et al [6], who use the basic agglomerate model for 21% and 10% oxygen feeds and get good agreement with experiment for both cases. Second, the fit for oxygen feed in the present work is actually closer to experimental data than the one predicted by Maya et al. [6]. However, because the estimation of mass-transport limited conditions is better for lean oxygen feed, all of the additional results generated in our study were based on air entering the gas channel. Model predictions of variations of platinum loading in the catalyst layer (obtained by changing the concentration of platinum in the catalyst, Pt/(Pt+C) ratio) were also compared with experimental measurements at constant electrode potential. The results are shown in Fig. 6 for the performance of Electrodes A in air and B in oxygen at potentials of 0.7 V and 0.6 V, the potentials at which the respective experimental data [20,27] were obtained. Both the modeling results and the experimental data suggest that at intermediate potentials of 0.6 -0.7 V, the increase in current density caused by increased platinum loading becomes more moderate as the platinum loading increases. The explanation for this trend is that, as the catalyst is made more concentrated in platinum, the size of the platinum particles (agglomerated crystallites) increases and the surface area of platinum that is available for reaction decreases. This means that part of the performance benefit associated with using more platinum is sacrificed and a linear relationship between current density and platinum loading, at constant potential, does not hold.

Figure 6. Effect of platinum loading on electrode performance. Electrode A – Air at 0.7 V. (experimental data [20] are shown as open squares, model predictions are shown as a solid line). Electrode B – Oxygen at 0.6 V. (experimental data [27] are shown as open circles, model predictions are shown as a dashed line).

Another way of changing the platinum content of the electrode is by changing the thickness of the catalyst layer for the same concentration of platinum [Pt / (Pt + C)] on carbon catalyst. The polarization curves predicted by the model for catalyst layers of different thickness are shown in Fig. 7 at conditions corresponding to the operation of Electrode B

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(Table 1 lists characteristics and conditions) in air. The platinum loading (mg Pt /cm2 CL) corresponding to each line is proportional to the catalyst layer thickness. The limiting current density (the current density at which the electrical potential becomes zero) is a useful criterion for comparing polarization curves. The limiting current density is predicted to increase substantially as the catalyst thickness is changed from 50 μm to 150 μm. The increase is considerably smaller as the thickness is changed from 150 μm to 250 μm and an appreciable decrease accompanies the increase in thickness from 250 μm to 450 μm. These results can be explained qualitatively. As the catalyst layer thickness increases at least two changes occur. More platinum reaction sites are added and the effect of ohmic polarization (transport of protons from the ELL to reaction sites in the CL) becomes larger. The addition of catalyst sites has the greatest impact on the thinnest catalyst layers. However, the contribution of ohmic polarization becomes progressively more dominant as the catalyst layer becomes thicker. Ohmic polarization also becomes progressively more dominant as the current density increases. The polarization curve for the catalyst thickness of 450 μm illustrates this point. At small current densities, it presents the largest electrode potential. However, at large current densities, a smaller limiting current density is predicted than two of the electrodes that have smaller catalyst loadings. When the catalyst layer becomes very thick, the incremental negative effect of the additional ohmic polarization is greater than the incremental benefit of additional platinum.

Figure 7. Model predictions of the effect of catalyst layer thickness (varying platinum loading) obtained at constant catalyst composition [ Pt/(Pt+C) ], on the current density for Electrode B in air. The catalyst thickness was 50, 100, 150, 250, and 450 μm.

The effect of phosphoric acid concentration on the performance of the oxygen electrode at constant temperature was also investigated. Polarization curves for Electrode A (Table 1) in air were predicted by the model for various phosphoric acid concentrations and are shown in Fig. 8. At 160°C the limiting current density is shown to decrease as the phosphoric acid concentration increases.

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Figure 8. Model predictions showing the effect of phosphoric acid concentration on the current density for Electrode A in air at constant temperature. Polarization curves are shown for H3PO4 concentrations of 85, 90, 100, and 105 %.

When the phosphoric acid concentration increases, the diffusivity of oxygen dissolved in the electrolyte decreases and the specific conductivity of the electrolyte also decreases causing both the ohmic polarization and the concentration polarization to increase. The net result is that the limiting current density decreases as the phosphoric acid concentration increases, as shown in Fig. 8. This trend may be reversed at smaller phosphoric acid concentrations because the phosphoric acid conductivity is smaller. The specific conductivity of H3PO4 solutions at constant temperature is maximized at intermediate concentrations (the value of the maximum depends on temperature). However, the correlation developed in this work is valid for acid concentrations greater than 75 wt% H3PO4 so no attempt was made to find an optimal acid concentration at a given temperature. Also, for small acid concentrations the permissible operating temperature becomes too low for practical applications, which is the main reason why concentrated electrolyte is normally used in PAFC fuel cell applications. The maximum permissible operating temperature is determined by the boiling point of the phosphoric acid, which in turn is a function of the phosphoric acid concentration. However, as the phosphoric acid concentration increases the water concentration decreases and therefore the water vapour partial pressure decreases. That means that solutions having larger H3PO4 concentrations have greater boiling points. To examine this effect, the temperature was increased while simultaneously increasing the phosphoric acid concentration. Increasing the temperature is expected to increase (a) the surface reaction processes, (b) the diffusivity of oxygen dissolved in the electrolyte, and (c) the specific conductivity of the acid, and therefore to decrease all types of polarization losses. The predicted performance of Electrode A in air is shown in Fig. 9 at four different combinations of temperature and phosphoric acid concentration (varying from 85 % H3PO4 / 155°C to 105 % H3PO4 / 240°C) for a constant water vapour pressure of 0.9 atm. The four lines in Fig. 9 are hardly distinguishable. The results from the model indicate that the beneficial effect of increasing the temperature essentially offsets the negative effect of increasing the phosphoric acid concentration.

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Figure 9. Model predictions of polarization curves for a cathode having the properties of Electrode A with air feed and with combinations of phosphoric acid concentration and temperature that result in the same water vapour pressure (0.9 atm.). Combinations of 85 % H3PO4 / 155°C, 90 % H3PO4 / 176°C, 100 % H3PO4 / 205°C, 105 % H3PO4 / 240°C, are represented by curves 1 to 4 respectively.

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The modeling results are in good agreement with the experimental data for operation in both oxygen and in air, as seen in Fig. 4 to 6. The polarization curves in Fig. 4 and 5 correspond to two electrodes having very different electrode characteristics, as seen in Table 1. This demonstrates that the present model can differentiate between the performance controlling factors for these two electrodes. It also converges and simulates the behavior of the polarization curve for close to limiting current density conditions for operation in air. It is interesting to note that both the model predictions and the experimental data indicate that Electrode A is performing better than Electrode B over the whole range of current densities even though they have comparable platinum loadings.

Figure 10. Model predictions of polarization curves for Electrode B with various resistances removed. Curve 1 – includes all resistances. Curve 2 – Removal of ionic conduction loss. Curve 3 – Removal of losses caused by both ionic conduction and diffusion of dissolved oxygen in the liquid in agglomerate pores. Curve 4 – Removal of losses from ionic conduction, diffusion of dissolved oxygen in the liquid in agglomerate pores, and diffusion of dissolved oxygen in the liquid film. Curve 5 – Removal of losses from ionic conduction, diffusion of dissolved oxygen in the liquid in agglomerate pores, diffusion of dissolved oxygen in the liquid film, and diffusion of gaseous oxygen in both the gas diffusion layer and the catalyst layer. Curve 6 – Ideal reversible potential at 190°C.

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There are several physical and chemical processes that occur at the electrode that cause the total losses in electrical potential to increase as the current density increases. These processes are the surface reaction processes (adsorption, electrochemical reaction, desorption), various molecular diffusion processes in the electrode, and ionic conduction in the catalyst layer. The relative contribution of each of these processes can be quantified using the mathematical model, in order to assess their relative influence on the cathode performance. The particular parameter in the model (electrolyte conductivity, dissolved oxygen diffusivity in the electrolyte, and gas phase diffusion coefficients) that was relevant to the diffusion and ohmic processes was individually forced to an extremely large value (l010) so that the relevant process was taken to occur extremely rapidly. By eliminating the resistance caused by each process, it was possible to assess the importance of that process on cathode performance. The results are shown in Fig. 10. Line (1) is the polarization curve with all of the processes included. The effect of ionic conduction was removed from Line (1) to obtain Line (2). As a result the difference between Line (1) and Line (2) represents the potential loss caused by ionic conduction. The diffusion of oxygen dissolved in the electrolyte in the agglomerate interior pores was removed from Line (2) to obtain Line (3). The diffusion of oxygen dissolved in the electrolyte film surrounding the agglomerates was removed from Line (3) to obtain Line (4). Finally the diffusion of oxygen in the gas filled pores within both the catalyst layer and the gas diffusion layer was removed from Line (4) to obtain Line (5). The horizontal line at 1.093 V represents the ideal reversible potential of the cathode at 190°C. The vertical distance between any two consecutive lines in Fig. 10 represents the loss in electrical potential that is caused by the process that was eliminated to obtain the lower line. The three arrow bars in Figure 10 represent the losses in electrical potential caused by the three major processes: the surface reaction processes on the electrode (surface polarization), the ionic conduction in the electrolyte within the catalyst layer (ohmic polarization or iR loss) and, the transport of both oxygen in the gas-filled pores and dissolved oxygen in the electrolyte (diffusion polarization). The diffusion of oxygen in the gas filled pores contributes very little of the loss in electrical potential. The increase in ionic conduction and diffusion losses at higher current-densities can be qualitatively explained on the basis of this model. As the current-density increases, the reaction rate increases, the rate of proton consumption also increases and, correspondingly, the ionic current- density also increases. Thus, a larger potential gradient in the electrolyte in the catalyst layer is required to support a larger reaction rate so that the ohmic losses increase. The increase in the potential losses resulting from oxygen diffusion in the agglomerates and their surrounding liquid film can also be qualitatively explained. Consumption of oxygen by the reaction inside an agglomerate creates an oxygen concentration distribution in the agglomerate and the liquid film. As the current density is increased, the concentration gradient becomes steeper close to the agglomerate's outside surface and a greater portion of the agglomerate is not well utilized because of the low concentration of dissolved reactant gas in the liquid electrolyte. For smaller dissolved oxygen concentrations, higher polarization (driving force) is required for the reaction to proceed at a specific current density, which results in smaller electrode potentials.

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Figure 11. Model predictions of polarization curves for varying agglomerate size in the cathode layer, Electrode potential (V) versus current density (A/cm2) for Electrode B. Curves are shown for electrodes composed of agglomerates whose radii vary from 0.01 to 5 μm.

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The effect of agglomerate radius on the shape of the polarization curve is shown in Fig. 11. As the agglomerates become smaller, the length of the diffusion path from the agglomerate exterior to its centre becomes shorter, thereby decreasing the diffusion resistances. The limiting current density in Figure 10 increases as the agglomerate radius decreases from 5 μm to 0.1 μm. However, it was predicted that there is very little change in performance when the agglomerate radius becomes smaller than 0.1 μm. The oxygen concentration profiles within the agglomerate help to explain this phenomenon.

Figure 12. Dissolved oxygen concentration profiles within agglomerates of varying size – Oxygen concentration (moles/cm3) versus position in the agglomerate (μm). Agglomerate radii vary from 0.5 to 3 μm.

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Oxygen concentration profiles are shown in Fig. 12 for agglomerates having different radii. For the 3 μm agglomerates almost all of the oxygen is consumed very close to the agglomerate edge (3 μm -2.5 μm). For 2 μm agglomerates there is a finite oxygen concentration over a greater proportion of the agglomerate radius before the concentration drops to essentially zero (2 μm – 1 μm). In contrast for the 1.0 μm and 0.5 μm agglomerates the oxygen concentration remains considerable in the whole thickness of the agglomerate. The oxygen concentration profiles for 0.1 and 0.01 μm are not shown in Fig. 12 because their exterior radii are so small that their dimensions are virtually at the origin of the radius scale used in Fig. 12. Nevertheless by extrapolating the trends in the profiles shown in Fig. 12, it is apparent that the oxygen concentration profiles for 0.1 and 0.01 μm agglomerates would become flatter as the agglomerate radius decreased. Uniform concentration profiles provide an explanation for the polarization curves in Fig. 11 being similar when the CL is composed of either 0.1 μm or 0.01 μm agglomerates. The effect of the local oxygen conversion was also predicted using the model. Each local oxygen conversion corresponds to a specific gas phase composition in the gas channel. For example, for air feed to the electrode, 0% conversion corresponds to 0.21 oxygen mole fraction in the gas channel of a fuel cell cathode, assuming no water vapor in the feed. Larger conversions correspond to oxygen mole fractions having values that depend on the amount of oxygen consumed and the amount of water in the vapour phase. The term local oxygen conversion is used to differentiate it from the overall conversion, which would be the local oxygen conversion at the gas channel exit. The one-dimensional model is capable of calculating the local current density at any cathode potential for any value of local oxygen conversion. As the conversion is increased along the gas channel of a fuel cell, the gas is progressively depleted of oxygen, so that the performance deteriorates. A two dimensional model would be required to describe fuel cell performance as a function of distance along a gas channel. Nevertheless, this one-dimensional model can describe fuel cell performance as a function of oxygen conversion along a gas channel.

Figure 13. Model predictions of contour lines for constant percent of electrical potential loss, as a function of both local oxygen conversion and of local current density (A/cm2). The numbers denote the loss in electrical potential (V).

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One way to illustrate the relationship described in the previous paragraph is shown in Fig.13 where the electrical potential loss at the cathode is plotted as a function of both local current density and local conversion. The potential loss (polarization) is defined as the difference between the equilibrium potential for oxygen reduction and the actual cathode potential. Fig. 13 presents the model-predicted contour lines of constant potential losses for electrode B with air feed, as a function of both local current density and local oxygen conversion. The numbers on each contour line show the potential loss in Volts. At any particular value of oxygen conversion, the potential losses increase as the current density increases because the various resistances increase, as discussed before for the zero conversion polarization curve (Fig. 10). Also, the potential losses increase as the conversion increases for a particular current density because the oxygen concentration (partial pressure) decreases. As the conversion is increased, there is less oxygen to react and the concentration distributions across the agglomerates, liquid film and gas-pores become less uniform, and less oxygen reaches the catalyst sites; a larger polarization is then required to support the reaction rate. For example, the limiting current density, corresponding to the line labeled EEQ, in Fig. 13 (for which the entire equilibrium electrode potential has been lost) decreases as the conversion increases, which is a significant decrease in performance at high reactant conversions. In other words, the maximum current density that can be achieved at a fuel cell cathode (at EEQ) is a function of the oxygen conversion. The converse is also true. Any pair of current density and conversion values is necessarily to the left of the EEQ line in Fig. 13. The polarization of the electrode can then be determined as a function of both variables, by referring to Fig. 13. Also, a cathode electrode would be expected to perform significantly worse at locations away from the cathode gas channel inlet where much of the oxygen has been converted by the reduction reaction.

Figure 14. Model predictions of contour lines for constant percent of electrical potential loss caused by cathode overpotential (surface polarization) as a function of both local oxygen conversion and local current density (A/cm2). The numbers denote the percent of total potential loss caused by surface polarization (%).

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In a fuel cell cathode, the net current density measured would be the integral of the local current densities from the inlet (zero conversion) to the exit of the oxygen channel. At steadystate operation, the cathode catalyst layer will be at constant electrical potential. In order to predict the net current density for a specific cathode electrode potential (or specific value of potential losses), the integration would have to be performed along one of the contour lines of Figure 13, from zero conversion to the conversion at the exit. These considerations lead to the question of how the extent of oxygen conversion affects each of the various resistances that cause potential losses. For this reason the model was used to examine the fraction of the total potential loss that was caused by each of the surface, diffusion and ionic conduction processes at different values of oxygen conversion and current-density. The results in Fig. 14-18 correspond to operation of electrode B in air. Fig.14 presents contour lines of constant percentage of potential loss contributed by the electrochemical reaction (activation or surface overpotential) as a function of local oxygen conversion and current density. The numbers on each contour line of the graph denote the percentage of the total loss contributed by the surface overpotential. A similar contour map was made for all other resistances: the ionic conduction potential loss (Fig.l5), the potential loss attributed to oxygen diffusion in the gas phase (Fig. 16), the potential loss attributed to oxygen diffusion in the liquid film (Fig. 17) and the potential loss attributed to oxygen diffusion in the agglomerates (Fig. 18). In all cases, the number label of each contour line denotes the percentage of the total potential loss contributed by the specific resistance described. The processes described by Fig. 16, 17 and 18 are all different modes of the electrode's concentration polarization. All of the lines in Fig. 14-18 are within the region to the left of the EEQ line in Fig.13, because any pair of values for the two axes variables has to be within this region, as discussed above. Thus the lines in Fig. 15-18 are interrupted at this boundary, which corresponds to the case where the potential of the electrode drops to zero. As an example of how these figures are read, note that from Fig. 13, the potential loss in the electrode at a gas channel location corresponding to 50% oxygen conversion and 0.4 A /cm2 current density, is 0.7 V. From Fig. 14, the percentage of total potential loss caused by surface polarization at these conditions is 66%. Thus, the potential loss caused by surface polarization at 50% oxygen conversion, 0.4 A / cm2, is approximately 0.66 x 0.7 = 0.462 Volts. Similarly the potential losses caused by all process can be calculated using Fig. 13-18. As shown in Figure 14, at low current densities the surface reaction contributes the greatest proportion of the loss. At any particular conversion, as the current density is increased all other processes (Fig. 15-18) become increasingly important. The effect of the current density on all the resistances was described before (Fig. 10). The same explanations are valid at any value of oxygen conversion. As the current density is increased at a constant value of oxygen conversion, (any horizontal cut in Fig. 15-18), the percentage of the total loss caused by each of the ionic conduction and diffusion processes increases.

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Figure 15. Model predictions of contour lines for constant percent of electrical potential loss caused by ionic conduction (ohmic polarization) as a function of both local oxygen conversion and local current density (A/cm2). The numbers denote the percent of total potential loss caused by ohmic polarization (%).

Figure 16. Model predictions of contour lines for constant percent of electrical potential loss caused by gas phase oxygen diffusion in both the gas diffusion layer and the catalyst layer as a function of both local oxygen conversion and local current density (A/cm2). The numbers denote the percent of total potential loss caused by gas diffusion in gas filled pores (%).

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Figure 17. Model predictions of contour lines for constant percent of electrical potential loss caused by dissolved oxygen in the liquid film surrounding the agglomerates as a function of both local oxygen conversion and local current density (A/cm2). The numbers denote the percent of total potential loss caused by dissolved gas diffusion in the liquid film (%).

Figure 18. Model predictions of contour lines for constant percent of electrical potential loss caused by dissolved oxygen diffusion in the agglomerate pores filled with liquid as a function of both local oxygen conversion and local current density (A/cm2). The numbers denote the percent of total potential loss caused by dissolved gas diffusion in the agglomerates (%).

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The effect of oxygen conversion can be explained with similar arguments. At any particular current density (any vertical cut in Fig. 16-18) as the oxygen conversion increases, all modes of concentration polarization increase. As the gas entering the electrode pores from the gas channel becomes less concentrated in oxygen most of the oxygen reacts at locations very close to the outside surface of the agglomerates. Thus, greater potential losses result (Fig 13) and the percentage of the total potential loss caused by diffusion polarization becomes greater (Fig 16-18). The percentage of the electrical loss contributed by ionic conduction is fairly independent of conversion (Figure 15). Ionic conduction is an IR-drop and therefore depends strongly on the current density (the small dependence on conversion is caused by current density distribution changes in the electrode at different conversions). As the mass transport resistances (concentration polarization) increase at high oxygen conversions, the percentage of the potential loss contributed by the surface polarization decreases (Fig. 14). In summary, as the oxygen conversion increases, the driving force for the reaction decreases, the surface reaction processes contribute a smaller percentage and the diffusion processes contribute a larger percentage of the total loss in electrical potential. The percentage of total loss attributed to diffusion in gas-filled pores is almost negligible at any conditions. The combination of Fig. 13 to 18 can be used to identify the dominant processes at any combination of local oxygen conversion and local current density. Two such combinations are shown in Table 2. Both are taken at 50% oxygen conversion. The combination in Table 2 that has a local current density of 0.1 A/cm2 would correspond to a smaller local power density and the local current density of 0.75 A/cm2 would correspond to a larger local power density. The distribution of the electrical losses is very different in these two examples. The magnitude of the power demand (or current density) has an important influence on the percentage of total electric loss attributed to each of the processes. An ideal cathode design would have to be optimized over all anticipated operating conditions along the gas channel, as well as over all anticipated net power demand conditions.

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TABLE 2 CONTRIBUTIONS TO LOSS OF ELECTRICAL POTENTIAL Oxygen conversion [%]

50

50

Current density [A/cm2]

0.1

0.75

Total loss in electrical potential (Fig 13) [V]

0.5

0.93

Potential delivered = 1.093 – loss [V]

0.593

0.163

Power density [W/cm2]

0.06

0.12

Reaction (Fig 14) [%]

82

53

Ionic conduction (Fig 15) [%]

4.5

19.6

Dissolved O2 – pore diffusion(Fig 18)[%]

12.5

20.9

Dissolved O2 – film diffusion(Fig 17)[%]

0.8

5.7

Gaseous O2 – pore diffusion (Fig 16) [%]

0.2

0.8

Percentage of electrical loss

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The electrical potential losses are affected by both the current density and by the oxygen conversion. Increasing the current and power densities at constant oxygen conversion, results in an increase of the potential loss that is caused by both ionic conduction and diffusion of dissolved O2. In contrast, increasing the oxygen conversion at constant current density also results in an increase of the potential loss which in this case is caused by diffusion processes whereas the potential loss caused by ionic conduction remains approximately constant. The entrance to the gas channel always corresponds to a condition of zero conversion and maximum local current density. In contrast, the exit of the gas channel corresponds to the maximum conversion (the total conversion) and minimum local current density. By following any line of constant electrical potential loss in Fig. 13, from the entrance to the exit, it is seen that by moving away from the gas channel entrance, the conversion increases and the current density decreases. Correspondingly, the contributions of the various resistances are different at different positions in the gas channel. Thus for any particular range of power densities the electrode should be optimized after consideration of the potential loss caused by the various resistances for the whole range of their variation along the gas-channel. Furthermore, if the power demand is decreased, then the net current density will decrease, and the electrode potential will increase (potential loss decreases). The electrical potential losses at each position in the gas channel will decrease, thus shifting the operating constant potential line of Fig. 13 to one of smaller electrical potential loss. Correspondingly, the distribution of the potential loss among the various resistances will be different. Thus, the selection of the electrode characteristics should be made after the entire range of total power densities for a fuel cell has been considered. The combinations of Fig 13-18 can be beneficial to the understanding of the variations in performance along the gas channels of fuel cells. A fuel cell's control system is usually designed to maintain the flow of air at a certain multiple of the stoichiometric air requirement. That multiple fixes the conversion at the exit of the cathode gas channel. The cathode potential is approximately constant for any particular power delivered from the unit cell and any particular average current density. Therefore the constant potential loss contour line in Fig. 13 is determined by the cell power demand and will extend from zero local oxygen conversion to the oxygen conversion specified by the multiple of the stoichiometric requirement. When the power demand is small, the electrical potential losses are small and the current density distribution along the cathode gas channel will be very uniform, as is shown in the line corresponding to 0.35 V potential loss in Fig. 13. In contrast, when the power demand is high, there will be a very non-uniform current density distribution along the cathode gas channel, as for example, for the line labeled 0.9 V potential loss in Fig. 13. In all cases, the current density will be greatest at the gas channel inlet and smallest at the gas channel exit. Fig. 13 can be used to identify local current densities at the inlet, at the exit and the variation in the current densities along the gas channel The combination of Fig. 14 to 18 can be used to identify the approximate contribution of each resistance mechanism at the inlet, at the exit and at different intermediate current densities along the gas channel. The values for the percentage of electrical loss in Table 2 suggest that some modification to the properties of Electrode B might be desirable. Altering the electrode properties to diminish the diffusion resistance of dissolved oxygen in the agglomerate pores would be beneficial. Also diminishing the loss caused by ionic conduction when operating at a current density of 0.75 A/cm2 would be helpful. In contrast not much would be gained by attempting to diminish ionic conduction at cathodes operating at a current density of 0.1 A/cm2. The

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main point is that a cathode having properties that are optimized to operate at one condition will not be operating optimally at all other conditions. Since the operating conditions at the gas channel exit of some fuel cells change considerably with variations in load, and since the operating conditions will vary along the length of the gas channel, the entire range of likely operating conditions should be taken into account when selecting the electrode properties. Even though the results presented in Fig. 13-18 apply specifically to the properties of Electrode B, they also provide a general indication of the combined influence of current density and oxygen conversion on fuel cell performance.

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CONCLUSIONS The model of a PAFC cathode described here successfully predicted experimentally measured polarization curves for two cathodes having very different characteristics. In addition the model prediction for the effect of platinum loading on current density was in good agreement with experimental data. The model was used to predict cathode performance as a function of operating conditions and macroscopic electrode design variables. The limiting current density was found to increase when the phosphoric acid concentration is decreased from 105 % to 85 %. This was related both to a decrease in the specific conductivity of the electrolyte and to a decrease in the diffusivity of oxygen dissolved in the electrolyte. However, an increase in the H3PO4 concentration permits the operating temperature to be increased. It was found that the enhancement in performance caused by increased temperature was almost exactly offset by the diminished performance from larger H3PO4 concentrations. One of the most powerful attributes of the model is its capacity to predict phenomena that occur on a scale of microscopic dimensions that is too small to permit experimental measurements. Typically catalyst layers are 1 mm thick or less and the agglomerate radii are of the order of a few micrometers. The model predicted that cathode performance would increase as a function of catalyst layer thickness from 50 μm to 250 μm. For the electrode modeled, it was found that a catalyst layer thickness greater than 250 μm would cause performance to decrease, because the benefit of additional platinum is surpassed by the performance loss caused by increased ionic polarization. The model predicted an improvement in cathode performance as the agglomerate radius is decreased from 5 μm to 0.1 μm. Further decreases in agglomerate radii have negligible effect on performance. Oxygen profiles inside agglomerates showed that the concentration of oxygen at the center of larger agglomerates is essentially zero resulting in underutilization of platinum, while the profiles become progressively more flat as the agglomerate size is decreased. The efficiency of a fuel cell can be related to the loss in electrical potential. Profiles of electrical potential loss were calculated as a function of both local current density (reaction rate) and local oxygen conversion. The variation in local current density along the gas channel as a function of local oxygen conversion was identified. The calculations identified the portion of that loss that was attributed to each of the following phenomena; cathode overpotential, ionic conduction, diffusion of oxygen gas, diffusion of oxygen dissolved in the liquid film around agglomerates in the catalyst layer, and in the agglomerate pore structure. One of the observations from these calculations is that the mechanism causing the major portion of the electrical loss may change along the gas channel as the oxygen conversion

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increases or it may change with changes in current density. Fuel cell design optimization should be done with respect to the total range of current densities and oxygen conversions that are expected during operation.

ACKNOWLEDGMENTS Financial assistance from the Natural Sciences and Engineering Research Council of Canada for the support of one of the authors, George Psofogiannakis, is gratefully acknowledged.

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LIST OF SYMBOLS cA gas phase concentration of component A [moles/cm3] cT total gas phase concentration [moles/cm3] concentration of oxygen dissolved in electrolyte [moles/cm3] cO2D D/IN concentration of dissolved oxygen at the interface between the agglomerate cO2 and the liquid film [moles/cm3] cAG concentration of component A in the gas phase [moles/cm3] c* solubility of oxygen in the liquid electrolyte [ (mole/cm3)LIQ / atmGAS ] c** solubility of oxygen in liquid electrolyte at one atmosphere of gas pressure [moles/(cm3)LIQ] D diffusivity of oxygen dissolved in liquid electrolyte [cm2/s] DE effective diffusivity of oxygen dissolved in liquid electrolyte within a porous material [cm2/s] DAB molecular diffusivity of component A in component B [cm2/s] effective molecular diffusivity of component A in component B in a porous DABE 2 material [cm /s] E electrode potential , φS – φL|Z=Z2 [V] EEQ reversible equilibrium potential [V] F Faraday constant, 96487 [Coulombs/equivalent] FA Fraction of pore volume in the CL that is occupied by liquid FA at which all interior pores of an agglomerate are filled with liquid FA,MAX fraction of CL pore volume that is within the interior of the agglomerates (FP)IN ΔG Α activation energy (difference in Gibbs energy between the reactant and the transition state) that is used for the exchange current density [Joules/mole] j current density [A/cm2 of platinum metal] j0 exchange current density [A/cm2 of platinum metal] j+ net ionic current density [A/cm2] k ionic conductivity of phosphoric acid electrolyte [S cm– 1] kE effective ionic conductivity of phosphoric acid electrolyte in electrode pores [S cm – 1 ] K Henry’s law constant LCL thickness of the catalyst layer [cm]

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LF thickness of the liquid film surrounding an agglomerate [cm] MA molecular weight of species A MPt Platinum loading per unit electrode surface area [g Pt/cm2] NA Flux of species A [moles/(cm2 s)] P Pressure [Pa] r radial coordinate within an agglomerate [cm] rAG radius of an agglomerate [cm] R gas constant, 8.3145 [Joule/(mole K)] RAG reaction rate of dissolved oxygen per agglomerate [moles/(s agglomerate) ] volume reaction rate of dissolved oxygen within an agglomerate [mole/(s RV-AG cm3AGGLOMERATE) ] RV volumetric reaction rate in the catalyst layer [ mole/(s cm3) SA platinum surface area per unit volume of agglomerate [cm2 Pt / cm3AGGLOMERATE) ] SPt surface area per unit mass of platinum [cm2 Pt/g Pt] T temperature [K] u fraction of agglomerates that are utilized (filled with electrolyte) VC volume fraction of the catalyst layer composed of carbon volume fraction of the catalyst layer composed of agglomerates VFAG VG volume fraction of gas in the catalyst layer VL volume fraction of liquid in the catalyst layer VLF volume fraction of the catalyst layer composed of liquid film VPt volume fraction of platinum in the catalyst layer VS volume fraction of solid in the catalyst layer volume fraction of pores within the agglomerates (agglomerate porosity) VP/AG W concentration of phosphoric acid [wt%] weight fraction of platinum in catalyst = mass of Pt / mass of (Pt+C) WF/Pt weight fraction of polytetrafluoroethylene in catalyst layer WPTFE z coordinate along catalyst layer (or gas diffusion layer) increasing towards the electrolyte side of electrode [cm] zE ion equivalents per mole [ion equivalent/ mole] αC cathodic transfer coefficient φS electrical potential in the solid phase [V] φL electrical potential in the liquid electrolyte phase [V] η overpotential [V] ρi density of species i [g/cm3] ρAG number of agglomerates per unit volume catalyst layer [agglomerates/ cm3) ρμAG number of utilized agglomerates per unit volume catalyst layer (agglomerates filled with liquid / cm3) (ΣV)A molecular diffusion volume for species A ψ parameter defined by Eq 11

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REFERENCES [1] [2]

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[3] [4] [5] [6] [7] [8] [9] [10] [11] [12] [13] [14] [15] [16] [17] [18] [19] [20] [21] [22] [23] [24] [25] [26] [27] [28] [29] [30] [31] [32] [33] [34] [35] [36] [37]

(a) Psofogiannakis, G.; Bourgault, Y.; Conway, B.E.; Ternan, M.; J. Appl. Electrochem. 2006, 36 115. (b) Cairns, E.J.; Adv. Electrochem. Electrochem. Eng. 1971, 8 , 337. Anahara, R.; in Fuel Cell Systems (eds.L.J.M.J. Blomen, M.N. Mugerwa), Plenum, New York, 1993, p.271. Newman, J.; Electrochim. Acta 1979, 214, 223. Vogel, W.; Lundquist, J.; Bradford, A.; Electrochim. Acta 1972, 17, 1735. Illev, I.; Gamburzev, S.; Kaisheva, A.; Mrha, J.; J. Appl. Electrochem. 1975, 5, 291. Maja, M.; Tosc, P.; Vanni, M.; J. Electrochem. Soc. 2001, 148, A1368. Newman, J.S.; Tobias, C.W.; 1962, 109, 1183. Brown, R.; Rockett, J.A.; J. Electrochem. Soc. 1966, 113, 865. Srinivasan, S,; Hurwitz, H.D.; Bockris, J.O’M.; J. Chem. Phys. 1967, 46, 3108. Horowitz, H.H.; J. Electrochem. Soc. 1967, 114, 650. Newman, J.S.; Tiedemann, W.; AIChE J. 1975, 21, 25. Will, F.G.; J Electrochem. Soc. 1963,110, 152. Austin, L.G.; Lerner, H.; Electrochim. Acta 1964, 9, 1469. Burshtein, R.C.; Markin, V.S.; Pshenichnikov, A.G.; Chismadgey, B.A.; Chirkov, Y.G.; Electrochim Acta 1964, 9, 773. Grens, E.A.; Ind. Eng. Chem. Fund. 1966, 5, 542. Austin, L.G.; Ariet, M.; Walker, R.D.; Wood, G.B.; Comyn, R.H.; Ind. Eng. Chem. Fund. 1965, 4, 321. Giner, J.; Hunter, C.; J. Electrochem. Soc. 1969, 116, 1124. Cutlip, M.B.; Electrochim. Acta 1975, 20, 767. Maggio, G.; J. Appl. Electrochem. 1999, 29, 171. Kunz, H.R.; Gruver, G.A.; J. Electrochem. Soc. 1975, 122, 1279. Iczkowski, R.P.; Cutlip, M.B.; J. Electrochem. Soc. 1980, 127, 1433. Ross, P.N.; J. Electrochem Soc. 1980, 127, 2655. Perry, M.L.; Newman, J.; Cairns, E.J.; J. Electrochem. Soc. 1998, 145, 5. Bjornborn, P.; Electrochim. Acta 1987, 32, 115. Yang, S.C.; Cutlip, M.B.; Stonehart, P.; Electrochim. Acta 1989, 34, 703. Yang, S.C.; Cutlip, M.B.; Stonehart, P.; Electrochim. Acta, 1990, 35, 869. Maoka, T.; Electrochim. Acta 1988, 33, 371. Yang, S.C.; J. Electrochim. Acta 1988, 33, 371. Choudhury, S.R.; Deshmukh, M.B.; Rengaswamy, R.; J. Power Sources 2002, 112, 137. Psofogiannakis, G.; M.A.Sc. Dissertation, University of Ottawa 2003. Giordano, N.; Passalacqua, E.; Alderuci, V.; Staiti, P.; Lino, P.; Mirsalian, H.; Taylor, E.J.; Wilemski, G.; Electrochim. Acta 1991, 36, 1049. Watanabe, M.; Sei, H.; Stonehart, P.; J. Electroanal. Chem. 1989, 261, 375. Fuller, T.F.; Luczak, F.J.; Wheeler, D.J.; J.Electrochem. Soc. 1975, 142, 1752. Bird, R.B.; Stewart, W.E.; Lightfoot, E.N.; Transport Phenomena, 2nd ed. Wiley, New York 2002, 570.

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[38] Rothfeld, L.B.; A.I.Ch.E. J. 1963, 9, 19. [39] Poling, B.E.; Prausnitz, J.M.; O’Connell, J.P.; The Properties of Gases and Liquids, 5th ed. [40] McGraw Hill, New York, 2001. [41] Bird, R.B.; Stewart, W.E.; Lightfoot, E.N.; Transport Phenomena, 2nd ed. Wiley, New York [42] 2002, 565. [43] Geankoplis, C.J.; Mass Transport Phenomena, Holt Rinehart Wilson, New York, 1972, 152. [44] Gubbins, K.E.; Walker, Jr., R.D.; J. Electrochem. Soc. 1965, 112, 469. [45] Scharifker, B.R.; Zelenary, P.; Bockris, J.O’M.; J. Electrochem. Soc. 1987, 134, 2714. [46] Gard, D.R.; in Kirk Othmer Encyclopedia of Chemical Technology 4th ed., (J.I. Kroschwitz, [47] M. Howe-Grant, eds.), vol. 18, Wiley, New York, 1991. [48] Appleby, A.J.; J. Electrochem. Soc. 1979, 117, 328. [49] Kunz, H.R.; Grover, G.A.; Electrochim. Acta 1978, 2, 219. [50] Dobos, D.; Electrochemical Data, Elsevier, New York, 1975, 43. [51] Burden, R.L.; Faires, J.D.; Numerical Analysis, 7th ed. Brooks-Cole, 2001, 653. [52] ibid, p.272.

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Chapter 9

ALCOHOL OXIDATION ON PT AND PD BASED ELECTROCATALYSTS Changwei Xu1 and Yexiang Tong2* School of Chemistry and Chemical Engineering, Guangzhou University, Guangzhou 510006, China1 MOE of the Key Laboratory of Bioinorganic and Synthetic Chemistry, School of Chemistry and Chemical Engineering, Institute of Optoelectronic and Functional Composite Materials, Sun Yat-Sen University, Guangzhou 510275, China2

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ABSTRACT Pt has been extensively investigated as the elctrocatalyst for methanol and ethanol electrooxidation in acid media. Some attention has paid to 1-propanol, 2-propanol and EG electrooxidation in acid media. The methanol gives the best performance and 2propanol has very high current at low potential on Pt in acid media. When the alcohol electrooxidation operates in alkaline instead of acidic media the activity for all alcohols oxidation is significantly improved. EG shows the best electrooxidation activity on Pt in alkaline media. Pd is not a good electrocatalyst for methanol and EG, but it shows excellently higher activity than Pt for 1-propanol, 2-propanol and ethanol electrooxidation in alkaline media. For decades, significant progress has been made in the development of other metal modified Pt and Pd electrocatalysts for alcohol oxidation. It is well known that the Pt alloys are Pt-Ru and Pt-Sn and the correlated ternary Pt-Ru-based and Pt-Sn-based catalysts. The PtNi and PdNi have attracted more attention for alcohol electrooxidation. Enhanced activity of alcohol electrooxidation such as lower onset potential and improved stability was responsible for the change in the electronic properties of Pt and Pd in PtNi and PdNi. We have dispersed the PtNi and PdNi nanoparticles on Ti foil with a porous structure which were successfully prepared by electrodeposition method. Aaddition of Ni significantly promotes catalytic activity of the Pt and Pd electrocatalysts for the methanol and ethanol electrooxidation. Oxide promoted Pt and Pd catalysts for methanol and ethanol electrooxidation have been extensively investigated. In the oxides, the CeO2 and WO3 have attracted enormous attention. We have studied the alcohol electrooxidation on Pt-oxide and Pd-oxide in alkaline media. The addition of oxide to Pd promotes significantly the catalytic activity

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Changwei Xu and Yexiang Tong for the ethanol oxidation. The onset potential shifts the negative direction and the current density increases for the ethanol oxidation reaction on Pd-oxide/C in comparison to that on Pd/C.

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1. INTRODUCTION Direct alcohol fuel cells (DAFCs) based on methanol as fuel have attracted enormous attention as power sources for portable electronic devices and transportation due to the much higher energy density than gaseous fuels such as hydrogen and natural gas [1-3]. However, the development of DAFCs based on methanol fuel is facing serious difficulties [4, 5]: (i) slow electro-kinetic of methanol oxidation, (ii) high methanol crossover and (iii) high toxicity of methanol. Therefore, other alcohols were considered as alternative fuels. Ethanol is less toxic compared to methanol, and can be easily produced in great quantity by fermentation of sugarcontaining raw material. Ethanol is also the major renewable biofuel from the fermentation of biomass. Thus, direct ethanol fuel cells (DEFCs) have attracted more and more attention [6, 7]. The saturated C3-alcohols (1-propanol and 2-propanol) were investigated as the fuels for DAFCs. Especially, direct 2-propanol fuel cells have attracted more and more attention as 2propanol is the smallest secondary alcohol, less toxic than methanol and its electrochemical oxidation is of great interest due to its particular molecular structure [8]. The direct alcohol fuel cells using 2-propanol as fuel show much higher performance than direct methanol fuel cells and a much lower crossover current [9,10]. Polyhydric alcohols such as ethylene glycol (EG) have high boiling points and are much less volatile. They have higher theoretical energy density than that of methanol and can be electrochemically oxidized [11, 12]. The application of EG electrooxidation on DAFCs has attracted increased interests [13-15]. The fuel cells based on EG fuel show certain advantages such as low fuel cross-over and high power density [16, 17]. Pt has been extensively investigated as the elctrocatalyst for alcohol electrooxidation in acid media [18]. However, Pt has a low activity for alcohol electrooxidation in acid media. If alcohol electrooxidation proceeds in alkaline instead of acidic media, the kinetics will be significantly improved [19-21]. A lot of work has been done to study the electrooxidation of alcohol on Pt-based catalysts in alkaline media. However, the high price and limited supply of Pt constitutes a major barrier to the development of DAFCs. Pt-free electrocatalysts may be used as electrocatalysts for the alcohol oxidation in alkaline media. Our previous work on the development of Pt-free electrocatalysts for alcohol oxidation has focused on Pd-based catalysts and the results revealed that Pd is a good electrocatalyst for alcohol oxidation in alkaline media [22-31]. One effective approach to the cost reduction is to reduce the usage of the Pt and Pd catalysts. For decades, significant progress has been made in the development of metal and oxide modified Pt and Pd electrocatalysts for alcohol oxidation. We have reported recently that oxide promoted Pt and Pd catalysts for alcohol oxidation are much more active than pure Pt and Pd [23, 25, 30, 32-35].

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2. ALCOHOL ELECTROOXIDATION ON PT AND PD

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Pt has been extensively investigated as the elctrocatalyst for methanol [36-41] and ethanol electrooxidation [42-43] in acid media. Some attention has paid to 1-propanol [44], 2propanol [45, 46] and EG [47-49] electrooxidation in acid media. Fig. 1 shows the cyclic voltammograms of methanol, ethanol, 1-propanol, 2-propanol and EG oxidation in 0.5 mol L-1 H2SO4 solution containing 1.0 mol L-1 alcohol on Pt electrode. The working electrodes were platinum and palladium disk (99.999%) with a geometrical area of 0.03 cm2. A platinum foil (3.0 cm2) and a saturated calomel electrode (SCE, 0.241 V versus RHE) were used as counter and reference electrodes, respectively. A salt bridge was used between the cell and the reference electrode. The sweep rate is 50 mV s−1 in the potential range from -0.22 to 0.9 V. By comparing to the cyclic voltammogram (CV) in the absence of alcohol, an alcohol oxidation peak can be clearly observed for all alcohols. The electrooxidation of alcohol on electrode was characterized by two well-defined current peaks on the forward and reverse scans. In the forward scan, the oxidation peak is corresponding to the oxidation of freshly chemisorbed species which come from alcohol adsorption. The reverse scan peak is primarily associated with removal of carbonaceous species which are not completely oxidized in the forward scan [50, 51]. The magnitude of the peak current on the forward scan indicates the electrocatalytic activity of the electrocatalysts for the oxidation reaction of alcohol in both acid and alkaline media.

Figure 1. CVs of alcohols oxidation on Pt electrode in 1.0 mol L-1 alcohol + 0.5 mol L-1 H2SO4 solution at a scan rate of 50 mV s-1, 298 K.

The electrochemical performances of the oxidation reaction of methanol, ethanol, 1propanol, 2-propanol and EG on Pt electrode in 0.5 mol L-1 H2SO4 solution containing 1.0 mol L-1 alcohol were given in Table 1. It is clear that the current densities of methanol oxidation at corresponding potentials are higher than that of ethanol, 1-propanol and EG oxidation on Pt electrode. However, the onset potential (Es) and peak potential (Ep) for 2propanol electrooxidation are 100 mV and 340 mV more negative than that for methanol electrooxidation. More than 100 mV reduction on the onset potential for anodic reaction is significant for liquid fuel cell. Further, the 2-propanol electrooxidation on Pt occurs at significantly higher current at low potential than methanol does in acid meida. The results

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show that methanol shows the highest electrooxidation activity and 2-propanol has very high current at low potential on Pt in acid media.

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Table 1. Electrochemical performances of methanol, ethanol, 1-propanol, 2-propanol and EG oxidation on Pt in 0.5 mol L-1 H2SO4 solution containing 1.0 mol L-1 alcohol Alcohol

Es / V

Ep / V

jp / mA cm-2

j at 0.4 V / mA cm-2

CH3OH

0.13

0.62

4.5

0.58

ethanol

0.12

0.66

1.6

0.31

1-propanol

0.18

0.73

0.6

0.15

2-propanol

0.03

0.28

0.8

EG

0.16

0.67

0.7

0.05

Pt has a low activity for alcohol electrooxidation in acid media. If alcohol electrooxidation proceeds in an alkaline instead of an acidic media, the kinetics will be significantly improved [52]. So, the electrooxidation of methanol [53], ethanol [54], 1propanol [55], 2-propanol [56] and EG [57] has been extensively investigated on Pt in alkaline media. Fig. 2 shows the CVs of methanol, ethanol, 1-propanol, 2-propanol and EG oxidation in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol on Pt electrode. The sweep rate is 50 mV s−1 in the potential range from -0.95 to 0.17 V. By comparing to the CV in the absence of alcohol, an alcohol oxidation peak can be clearly observed for all alcohols.

Figure 2. CVs of alcohols oxidation on Pt electrode in 1.0 mol L-1 alcohol + 1.0 mol L-1 KOH solution at a scan rate of 50 mV s-1, 298 K.

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The electrochemical performances of the oxidation reaction of methanol, ethanol, 1propanol, 2-propanol and EG in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol on Pt electrode were given in Table 2. It is clear that the current densities of EG oxidation at corresponding potentials are higher than that of methanol, ethanol, 1-propanol and 2-propanol oxidation on Pt electrode. The oxidation reaction of 1-propanol and 2-propanol shows a very low activity on Pt electrode in alkaline media. The current density at -0.4 V of methanol oxidation is 2.9 times and 5.5 times than that of 1-propanol and 2-propanol oxidation. The activity order of alcohol oxidation on Pt electrode is EG > methanol > ethanol > 1-propanol > 2-propanol. Table 2. Electrochemical performances of methanol, ethanol, 1-propanol, 2-propanol and EG oxidation on Pt and Pd electrodes in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol Es / V

Ep / V

jp / mA cm-2

j at -0.4 V / mA cm-2

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Alcohol Pt

Pd

Pt

Pd

Pt

Pd

Pt

Pd

CH3OH

-0.64

-0.46

-0.29

-0.24

77

87

31.2

4.8

ethanol

-0.73

-0.67

-0.32

-0.25

35

100

29.0

34.6

1-propanol

-0.66

-0.70

-0.29

-0.29

15

78

10.7

39.2

2-propanol

-0.74

-0.73

-0.32

-0.36

6

46

5.7

41.1

EG

-0.66

-0.46

-0.19

-0.18

229

173

36.2

2.5

However, the high price and limited supply of Pt constitutes a major barrier to the development of DAFCs. Pt-free electrocatalysts such as Au [58-60], Ag [61, 62], Ni [63-65] and Cu [66] may be used as electrocatalysts for the alcohol oxidation in alkaline media. Our previous work on the development of Pt-free electrocatalysts for alcohol oxidation has focused on Pd-based catalysts and the results revealed that Pd is a good electrocatalyst for alcohoil oxidation in alkaline media. The methanol [67] and ethanol [68, 69] electrooxidation on Pd have been studied by other group in alkaline media. Fig. 3 shows the CVs of methanol, ethanol, 1-propanol, 2-propanol and EG oxidation in 0.5 mol L-1 H2SO4 solution containing 1.0 mol L-1 alcohol on Pd electrode. The sweep rate is 50 mV s−1 in the potential range from -0.22 to 0.9 V. By comparing to the CV in the absence of alcohol, no any alcohol oxidation peak can be clearly observed in 0.5 mol L-1 H2SO4 solution containing 1.0 mol L-1 alcohol. The results show that Pd has no activity for alcohol oxidation in acidic media. Fig. 4 shows the CVs of methanol, ethanol, 1-propanol, 2-propanol and EG oxidation in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol on Pd electrode. The sweep rate is 50 mV s−1 in the potential range from -0.95 to 0.17 V. By comparing to the CV in the absence of alcohol, an alcohol oxidation peak can be clearly observed for all alcohols.

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Figure 3. CVs of alcohols oxidation on Pd electrode in 1.0 mol L-1 alcohol + 0.5 mol L-1 H2SO4 solution at a scan rate of 50 mV s-1, 298 K.

The electrochemical performances of the oxidation reaction of methanol, ethanol, 1propanol, 2-propanol and EG in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol on Pd electrode were given in Table 2. Methanol and EG both show higher electrooxidation activity than ethanol, 1-propanol and 2-propanol in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol on Pt electrode. However, the 1-propanol and 2-propanol which both show very low activity on Pt electrode give good electrochemical performances on Pd electrode in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 alcohol. The current densities at -0.4 V of 1propanol and 2-propanol oxidation are 8.2 times and 8.6 times than that of methanol oxidation on Pd electrode and higher than that of ethanol oxidation. The onset potentials (Es) of 1-propanol and 2-propanol oxidation are more negative about 240 mV and 270 mV compared to that of methanol oxidation on Pd electrode and more negative than that of ethanol oxidation. The peak currents of 1-propanol and 2-propanol oxidation on Pd electrode begin to rise much more sharply at more negative potential, which will directly improve the fuel cell efficiency. The results show that 1-propanol and 2-propanol oxidation has much higher catalytic activity than methanol electrooxidation on Pd electrode in alkaline media. The activity order of alcohol oxidation on Pd electrode is 1-propanol ≈ 2-propanol > ethanol > methanol > EG. On the other hand, the current densities at -0.4 V for 1-propanol and 2propanol oxidation on Pd electrode are about 3.7 times and 7.2 times than that on Pt electrode, respectively. Pd shows better performances than Pt for 1-propanol, 2-propanol and ethanol oxidation in alkaline media.

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Figure 4. CVs of alcohols oxidation on Pd electrode in 1.0 mol L-1 alcohol + 1.0 mol L-1 KOH solution at a scan rate of 50 mV s-1, 298 K.

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3 ALCOHOL ELECTROOXIDATION ON PT-M AND PD-M One effective approach to the cost reduction is to reduce the usage of the Pt and Pd catalysts. On the other hand, the partial oxidation intermediates such as -CH3O, -CH2O, CHO, and -CO, which adsorb strongly on the surface of the catalysts resulting in deactivation [70-72]. The partial oxidation intermediates can adsorb on the electrodes which can block sites for alcohol adsorption. For example, compared to the lowest potentials that enable the formation of CO from methanol on Pt, potentials at least 0.1~0.2 V more positive are required to activate water and produce the surface oxides needed to convert CO to CO2 [73]. For decades, significant progress has been made in the development of other metal modified Pt and Pd electrocatalysts for alcohol oxidation. It is well known that the Pt alloys are Pt-Ru and Pt-Sn and the correlated ternary Pt-Ru-based and Pt-Sn-based catalysts. PtRu shows better performances than Pt for methanol electrooxidation both in acid [74-80] and alkaline [20] media. The other alcohol electrooxidation on PtRu [14, 81, 82] and PdRu [83] has also been studied. The alcohol electrooxidation on Pt-Ru catalyst is explained as a bi-functional mechanism [84]. Pt acts as main catalyst for catalysing the dehydrogenation of alcohol during the oxidation reaction and oxygen-containing species (OHad) can form on the Ru surface at lower potentials. These oxygen-containing species react with CO-like intermediate species on the Pt surface to produce RCO2 and release the active sites [85]. Ru-OH2 → Ru-OH + H+ + e

(1)

Pt-RCO + Ru-OH → Pt + Ru + RCO2 + H+ +e

(2)

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The PtSn is a very active catalyst for ethanol eletrooxidation in acid meida [86-88]. The mechanism was pointed out that the inhibition of C-C bond cleavage reactions by addition Sn to Pt catalysts was been attributed to changes in the geometric and electronic characteristics of the surface [7, 89]. The formation of CO2 and acetaldehyde is lowered whereas the formation of acetic acid is increased by addition of tin to platinum [90]. The other Pt alloys such as PtSn [91, 92], PtTi [93], PtBi [94, 95], PtOs [96, 97], PtAu [98-100], PtPd [101, 102], PtCo [103], PtRh [104, 105], PtMo [106] have been extensively investigated as alcohol electrooxidation catalysts in acid and alkaline media. The PtNi [107-110] and PdNi [111] have attracted more attention for alcohol electrooxidation. Some researchers found that PtNi nanoparticles showed a shift in the Pt 4f peak in a PtNi-based alloy structure, so called the electronic effect [108, 112]. Enhanced activity of alcohol electrooxidation such as lower onset potential and improved stability was responsible for the change in the electronic properties of Pt and Pd in PtNi and PdNi. In the PtNi and PdNi alloy nanoparticles, the Ni states on the Ni surface consist of Ni metal as well as Ni oxide and hydroxides (NiO, Ni(OH)2, and NiOOH) [112]. It has recently been reported that nickel or nickel hydroxides are not only promoters but also catalysts capable of oxidizing methanol in alkaline media [113, 114]. The following reaction schemes were nickel hydroxide as the catalyst for methanol oxidation: 2Ni3+O2- + CH3OH →Ni2+OH- + Ni2+O-CH2OH →Ni3+O2- + Ni3+O-C-HO + 3H+ +e Ni3+O-C-HO → Ni3+O-C-=O + H+ + e

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→ Ni3+O2- + CO2 + 2H+ + 2e

(3) (4) (5) (6)

The alcohol oxidation mechanism at the Ni electrode in alkaline media has been proposed. The proposed mechanism may be summarized as follow [115, 116] OH- + Ni(OH)2 ↔ NiOOH + H2O + e

(7)

NiOOH + RCH2OH → Intermediate 1 + Ni(OH)2

(8)

NiOOH+ Intermediate 1 → RCHO + Ni(OH)2

(9)

NiOOH + RCHO → Intermediate 2 + Ni(OH)2

(10)

NiOOH + Intermediate 2 → RCOOH + Ni(OH)2

(11)

Here, Pt-Ni and Pd-Ni porous films were prepared by electrodeposition method and were used as catalysts for ethanol electrooxidation in alkaline media. The Pt-Ni alloy film was deposited on Ti by 1 mA cm-2 electrodeposition in 0.01 mol L-1 H2PtCl6 + 0.007 mol L-1 NiCl2 + 0.1 mol L-1 KCl + 0.5 mol L-1 citric acid and control time to get 0.1 C cm-2. Fig.5a shows the TEM image of Pt-Ni nanoparticles deposited on Ti. The distribution of

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nanoparticles appears to be uniform with a slice form and a porous structure. The thickness of slice is about 12 nm. The results indicate that the Pt nanoparticles on Ti were successfully prepared by the electrodepostion method used in this study. The composition was confirmed by energy dispersive X-ray spectrometer (EDS) analysis and the results show that Pt and Ni elements are 51.4 and 48.6 by atomic content in the thin film. The loading of Pt and Ni is 0.032 and 0.011 mg cm-2.

Figure 5. SEM images of (a) the PtNi alloy film prepared by 1 mA cm-2 electrodeposition in 0.01 mol L-1 H2PtCl6 + 0.005 mol L-1 NiCl2 + 0.1 mol L-1 KCl + 0.5 mol L-1 citric acid; (b) PdNi alloy film prepared by 2 mA cm-2 electrodeposition in 0.01 mol L-1 PdCl2 + 0.05 mol L-1 NiCl2 + 0.1 mol L-1 KCl +0.5 mol L-1 citric acid, 0.1 C cm-2.

The Pd-Ni alloy film was deposited on Ti by 2 mA cm-2 electrodeposition in 0.01 mol L-1 PdCl2 + 0.05 mol L-1 NiCl2 + 0.1 mol L-1 KCl + 0.5 mol L-1 citric acid and control time to get 0.1 C cm-2. Fig.5b shows the TEM image of Pd-Ni nanoparticles deposited on Ti. The distribution of particles appears to be uniform with a flower form and a porous structure. The diameter of particles is about 50-60 nm. The composition was confirmed by EDS analysis and the results show that Pd and Ni elements are 78.4 and 21.6 by atomic content in the thin film. The loading of Pd and Ni is 0.043 and 0.007 mg cm-2.

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Fig. 6 shows the CVs of the ethanol oxidation on Pt (Pt loading: 0.051 mg cm-2) , Pd (Pd loading: 0.055 mg cm-2) , PtNi (Pt loading: 0.032 and Ni loading: 0.011 mg cm-2) and PdNi (Pd loading: 0.043 and Ni loading: 0.007 mg cm-2) in a N2-saturated (a) 1.0 mol L-1 KOH and (b) 1.0 mol L-1 KOH + 1.0 mol L-1 ethanol solution at a sweep rate of 50 mV s-1. The area of H-desorption deducted the double layer charging on the cyclic voltammogram curves in 1.0 mol L-1 KOH represents the charge passed for H-desorption QH, and can be used to estimate the electrochemical active surface (EAA) of the electrode [117]. The EAAs are 7 m2 g-1 for Pt, 6 m2 g-1 for Pt-Ni, 6 m2 g-1 for Pd and 4 m2 g-1 for Pd-Ni. The results show that the EAAs of PtNi and PdNi are not higher than that of Pt and Pd, indicating that the ethanol electrochemical oxidation is more active on PtNi and PdNi electrode than on Pt and Pd electrodes without increasing the EAA of the electrode. The decreasing of EAAs for Pt and Pd by adding Ni is due to the Ni atoms take the places of Pt and Pd atoms which are the electrochemical active centers.

Figure 6. CVs of the ethanol oxidation on Pt (Pt loading: 0.051 mg cm-2) , Pd (Pd loading: 0.055 mg cm-2) , PtNi (Pt loading: 0.032 and Ni loading: 0.011 mg cm-2) and PdNi (Pd loading: 0.043 and Ni loading: 0.007 mg cm-2) in a N2-saturated (a) 1.0 mol L-1 KOH and (b) 1.0 mol L-1 KOH + 1.0 mol L-1 ethanol solution at a sweep rate of 50 mV s-1.

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The onset potential of ethanol oxidation is -0.54 V on Pt electrode and it negatively shifts to -0.74 V on PtNi electrode. Similar results are observed for ethanol oxidation on Pd and PdNi electrodes. The onset potential for ethanol oxidation is -0.61 V on Pd electrode and it negatively shifts to -0.71 V on PdNi electrode. More than 100 mV reduction on the onset potential for anodic reaction is significant for liquid fuel cell. The change in the onset potential shows an improvement in the kinetics due to the synergistic effect by the interaction between Pt or Pd and Ni. As shown in the Fig. 6, the current density for ethanol oxidation at the potential of -0.4 V on Pt-Ni electrode is 33.1 mA cm-2 and higher than 8.8 mA cm-2 on Pt. Similar results are observed for ethanol oxidation on Pd and PdNi electrodes. The current density for ethanol oxidation at the potential of -0.4 V on PdNi electrode is 39.0 mA cm-2 and higher than 21.1 mA cm-2 on Pd. The activity for ethanol electrooxidation on Pt and Pd electrodes is significantly increased by the addition of Ni.

Figure 7. CVs of the ethanol oxidation on Pd (Pd loading: 0.3 mg cm-2) , PdCeO2 (Pd loading: 0.3 mg cm-2 and CeO2 loading: 0.015 mg cm-2) , PdNiO (Pd loading: 0.3 mg cm-2 and NiO loading: 0.005 mg cm-2), PdCo3O4 (Pd loading: 0.3 mg cm-2 and Co3O4 loading: 0.008 mg cm-2) and PdMn3O4 (Pd loading: 0.3 mg cm-2 and Mn3O4 loading: 0.008 mg cm-2)in a N2-saturated 1.0 mol L-1 KOH + 1.0 mol L-1 ethanol solution at a sweep rate of 50 mV s-1.

4. ALCOHOL ELECTROOXIDATION ON PT-OXIDE AND PD-OXIDE Research revealed that hydrous RuO2 was a more active catalyst for alcohol oxidation than did Ru0 as part of bimetallic Pt–Ru alloy [118-123]. The results can be interpreted as being associated with the presence of hydrous RuO2 (RuO2-δ(OH)δ), which would play the role of donor of the oxygen-containing species that promote the CO to CO2 oxidation. RuO2 is rapidly and reversibly oxidized and reduced by electrochemical protonation [124]. RuO2 + δ H+ + δ e ↔ RuO2-δ (OH)δ 0 ≤ δ ≤ 2

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(12)

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A lot of papers have been devoted to the effect of oxide promoted Pt and Pd for methanol and ethanol electrooxidation on Pt and Pd catalysts. In the oxides, the CeO2 has attracted enormous attention. The PtCeO2 base catalysts have been studied for methanol [125-129] and ethanol [130] electrooxidation in acid media. A mechanism for ceria’s proposed enhancement of alcohol oxidation can be described by the following reactions [131]: CH3OH + H2O ↔ Pt-COads + 4H+ + 4e

(13)

2CeO2 + 2H+ + 2e ↔ Ce2O3 + 2H2O

(14)

Pt + COads + 2CeO2 + 2H+ 2e ↔ Ce2O3 + Pt + CO2 + H2O

(15)

RCH2OH→ Pt-RHCOads + 2H+ +2e

(16)

Pt + RHCOads + 2CeO2 + 2H+ +2e ↔ Ce2O3 + Pt + RCOOH + H2O

(17)

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The PtWO3 has also attracted enormous attention for methanol [132, 133] and ethanol [134] in acid media. The WO3 improvement Pt activity for methanol oxidation is partly attributed to the ability of WO3 to intercalate H atoms into its matrix, forming tungsten bronzes. This feature of WO3 allows for the spillover of hydrogen from Pt sites into the WO3 matrix [135]. (18) WO3 + xPt-H → HxWO3 + xe + x Pt → WO3 +xe + xH+ + xPt The transfer of hydrogen ions, produced on the platinum during the electrooxidation of methanol, to the tungsten oxide ensures that the active reaction sites on the platinum remain clean, thus enhancing the electrooxidation current density. The other Pt-oxide and Pd-oxide such as Pt-TiO2[137], Pt-MoO3[138, 139], PtSnOx[140], Pt-LnOx(Ln = Sc, Y, La, Ce, Pr and Nd)[141], Pt-Fe2O3[142], PtZrO2[143], PtMgO[35, 144] and PdNiO[145] have been extensively investigated as alcohol electrooxidation in acid and alkaline media. We have studied the alcohol electrooxidation on Pt-CeO2 [32-34], PtNiO [23, 30], PtMgO [35], PdCeO2 [23, 25], PdNiO [23, 25, 29], PdCo3O4 [25] and PdMn3O4 [25]in alkaline media. Pt/C and Pd/C promoted with the oxides give a much higher activity and significantly enhanced performance stability for the electrooxidation reactions of alcohol than that on unpromoted Pt/C and Pd/C in alkaline media. Fig. 7 shows the CVs of the ethanol oxidation reaction on Pt/C, Pd/C, Pt-CeO2(2:1, w:w)/C, Pd-NiO(6:1, w:w)/C, Pd-Co3O4(4:1, w:w)/C and Pd-Mn3O4(4:1, w:w)/C electrodes in a N2-saturated 1.0 mol L-1 KOH solution containing 1.0 mol L-1 ethanol at a sweep rate of 50 mV s-1. The electrocatalyst powders were dispersed in 2-propanol with 5 wt% Nafion solution under ultrasonic stirring and then the catalyst ink was deposited on the surface of a graphite rod with a geometric area of 0.50 cm2 and dried at 80oC for 30 minutes. The platinum and palladium loading on the electrodes was normally controlled at 0.3 mg cm-2. An anodic peak current for the ethanol oxidation reaction can be clearly observed on all electrocatalysts. The onset potential, peak potential and current densities at peak potential and

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at the potential of -0.4 V are given in Table 3. The addition of oxide to Pd/C promotes significantly the catalytic activity for the ethanol oxidation. The onset potential shifts in the negative direction and the current density increases for the ethanol oxidation reaction on Pdoxide/C in comparison to that on Pd/C. Es of the ethanol oxidation on Pd-NiO(6:1, w:w)/C is 70 mV more negative than that on Pd/C. The results show that Pd-oxide/C has significantly higher activity for ethanol electrooxidation than Pd/C. Table 3: Onset potential (Es), peak potential (Ep), current density at Ep and at -0.3 V of Pt/C, Pd/C and Pt-oxide/C electrocatalysts for ethanol oxidation reaction in 1.0 mol L-1 KOH solution containing 1.0 mol L-1 ethanol at a sweep rate of 50 mV s-1. Pt or Pd loading: 0.3 mg cm-2. Electrocatalyst Pd/C Pd-CeO2(2:1, w:w)/C Pd-NiO(6:1, w:w)/C Pd-Co3O4(4:1, w:w)/C Pd-Mn3O4(4:1, w:w)/C

Es / V -0.66 -0.71 -0.73 -0.70 -0.67

Ep / V -0.19 -0.20 -0.19 -0.25 -0.19

jp / mA cm-2 19 39 95 47 57

j at -0.4V / mA cm-2 11.8 28.3 15.1 40.4 13.3

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5. CONCLUSIONS Pt has been extensively investigated as the elctrocatalyst for methanol and ethanol electrooxidation in acid media. Some attention has paid to 1-propanol, 2-propanol and EG electrooxidation in acid media. The methanol gives the best performance and 2-propanol has very high current at low potential on Pt in acid media. When the alcohol electrooxidation operates in alkaline instead of acidic media the activity for all alcohols oxidation is significantly improved. The activity order of alcohol oxidation on Pt is EG > methanol > ethanol > 1-propanol > 2-propanol. EG shows the best electrooxidation activity on Pt in alkaline media. The results show that Pd has no activity for all alcohols oxidation in acidic media, however Pd is a good electrocatalyst for alcohol oxidation in alkaline media. All alcohols show a considerable activity. The activity order of alcohol oxidation on Pd is 1propanol ≈ 2-propanol > ethanol > methanol > EG. Pd is not a good electrocatalyst for methanol and EG, but it shows excellently higher activity than Pt for 1-propanol, 2-propanol and ethanol electrooxidation in alkaline media. For decades, significant progress has been made in the development of other metal modified Pt and Pd electrocatalysts for alcohol oxidation. It is well known that the Pt alloys are Pt-Ru and Pt-Sn and the correlated ternary Pt-Ru-based and Pt-Sn-based catalysts. The PtNi and PdNi have attracted more attention for alcohol electrooxidation. Some researchers found that Pt/Ni nanoparticles showed a shift in the Pt 4f peak in PtNi-based alloy structure, so called the electronic effect. Enhanced activity of alcohol electrooxidation such as lower onset potential and improved stability was responsible for the change in the electronic properties of Pt and Pd in PtNi and PdNi. In the PtNi and PdNi alloy nanoparticles, the Ni states on the Ni surface consist of Ni metal as well as Ni oxide and hydroxides (NiO, Ni(OH)2, and NiOOH). It has recently been reported that nickel or nickel hydroxides are not

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only promoters but also catalysts capable of oxidizing methanol in alkaline media. We dispersed the PtNi and PdNi nanoparticles on Ti foil with a porous structure which were successfully prepared by electrodeposition method. Aaddition of Ni significantly promotes catalytic activity of the Pt and Pd electrocatalysts for the ethanol electrooxidation. A lot of papers have been devoted to the effect of oxide promoted Pt and Pd for methanol and ethanol electrooxidation on Pt and Pd catalysts. In the oxides, the CeO2 and WO3 have attracted enormous attention. We have studied the alcohol electrooxidation on Pt-oxide and Pd-oxide in alkaline media. The addition of oxide to Pt and Pd promotes significantly the catalytic activity for the ethanol oxidation. The onset potential shifts the negative direction and the current density increases for the ethanol oxidation reaction on Pd-oxide/C in comparison to that on Pd/C.

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REFERENCES [1] Varcoe, J. R.; Slade R. C. T. Fuel Cells 2005, 5, 187. [2] Arico, A. S.; Srinivasan, S.; Antonucci, V. Fuel Cells 2001, 1, 133. [3] Dyer, C. K. J. Power Sources 2003, 106, 31. [4] Heinzel, A.; Barragan, V. M. J. Power Sources 1999, 84, 70. [5] Ren, X.; Springer, T. E.; Zawodzmski, T. A.; Gottesfeld, S. J. Electrochem. Soc. 2000, 147, 466. [6] Song, S. Q; Tsiakaras, P. Appl. Catal. B 2006, 63, 187. [7] Antolini, E. J. Power Sources 2007, 170, 1. [8] Sumodjo, P. T. A.; Silva, E. J.; Rabochai, T. J. Electroanal. Chem. 1989, 271, 305. [9] Cao, D.X.; Bergens S. H.; J. Power Sources 2003, 124, 12. [10] Wang, J. T.; Wasmus, S.; Savinell, R.F. J. Electrochem. Soc. 1995, 142, 4218. [11] Belgsir, E. M.; Bouhier, E.; Yei, H. E.; Kokoh, H. K.; Beden, B. Electrochim. Acta 1991, 36, 1157. [12] Lebedeva, N. P.; Kryukova, G. N.; Tsybulya, S. V.; Salanov, A. N. Electrochim. Acta 1998, 44, 1431. [13] Matsuoka, K.; Inaba, M.; Iriyama, Y.; Abe, T.; Ogumi, Z.; Matsuoka, M. Fuel Cells 2002, 2, 35. [14] Wang, H.; Zhao, Y.; Jusys, Z.; Behm, R. J. J. Power Sources 2006, 155, 33. [15] Venancio, E. C.; Napporn, W. T.; Motheo, A. J. Electrochim. Acta 2002, 47, 1495. [16] Peled, E.; Livshits, V.; Duvdevani, T. J. Power Sources 2002, 106, 245. [17] Livshits, V.; Peled, E. J. Power Sources 2006, 161, 1187. [18] Steele, B. C. H.; Heinzel, A. Nature 2001, 414, 345 [19] Matsuoka, K.; Iriyama, Y.; Abea, T.; Matsuoka, M.; Ogumia, Z. J. Power Sources 2005, 150, 27. [20] Tripković, A. V.; Popović, K. D.; Grgur, B. N.; Blizanac, B.; Ross, P. N.; Marković, N. M. Electrochim. Acta 2002, 47, 3707. [21] Zhang, X.; Tsang, K. Y.; Chan, K.Y. J. Electroanal. Chem. 2004, 573, 1. [22] Xu, C. W.; Wang, H.; Shen, P. K.; Jiang, S. P. Adv. Mater. 2007, 19, 4256. [23] Xu, C. W.; Shen, P. K.; Liu, Y. L. J. Power Sources 2007, 164, 527. [24] Xu, C. W.; Cheng, L. Q.; Shen, P. K.; Liu, Y. L. Electrochem. Commun. 2007, 9, 997.

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Chapter 10

PHYSICOCHEMICAL AND ELECTROCATALYTIC PROPERTIES OF PtRu/C PREPARED BY IMPREGNATION REDUCTION METHOD: EFFECT OF PREPARATION PARAMETERS Jun Li2*, Jing-Dong Lin1and Dai-Wei Liao1* The State Key Laboratory of Physical Chemistry on Solid Surfaces, Department of Chemistry, College of Chemistry and Chemical Engineering, Institute of Physical Chemistry, Xiamen University, Xiamen 361005, China1 College of Power Engineering, Chongqing University, Chongqing 400030, China2

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ABSTRACT The impregnation reduction method was used to investigate the influence of preparation parameters on the physicochemical and electrocatalytic properties of 40 wt% PtRu/C. Based on the experimental results, it is found that the reductants with higher reducing power, such as formaldehyde, KBH4, methanol, N2H4, lead to smaller particles and higher dispersion. On the other hand, the reductants with lower reducing power, such as ethanol, isopropyl alcohol, result in bigger particles and poorer dispersion. In addition, the effect of pH value during impregnation was also investigated. In the case of using NaBH4 as a reductant, the favorable pH value is below 4, whereas in the case of using methanol, N2H4 and formaldehyde as the reductant, the favorable pH value is higher than 8. The characterization of the 40 wt% PtRu/C catalysts with X-ray diffraction (XRD) and transmission electron microscope show that the catalysts prepared under the condition of proper pH and α values (stoichiometric ratio of reductant to Pt and Ru metals) have welldispersed PtRu particles with average diameter around 2.3 nm. XRD experiments also indicate that the PtRu/C catalyst prepared by the reductant with different stoichiometric ratios and different pH values has different degrees of PtRu alloying. The effects of pH of precursor solution and α values of reductant on the average particle size of PtRu/C are discussed in detail. The electrocatalytic activity of the as-prepared catalysts was tested by *Corresponding authors at: Department of Chemistry, Institute of Physical Chemistry, Xiamen University, Xiamen 361005, China. Tel.: +86 592 218 3045 (D. W. Liao); College of Power engineering, Chongqing University, Chongqing 400030, China. Tel & Fax.:+86 23 6510 2474 (J. Li). E-mail addresses: [email protected] (D.W. Liao); [email protected] (J. Li).

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Jun Li, Jing-Dong Lin and Dai-Wei Liao cyclic voltammograms both in 1 M H2SO4 and 0.5 M H2SO4+ 1 M CH3OH. The results indicate that a reductant with higher reducing power, a high stoichiometric ratio of the reductant and a proper pH value are favorable to synthesize a catalyst with a higher electrocatalytic activity toward methanol oxidation. In addition, a new and highly reproducible method based on the flow injection (FI) technology was also developed to synthesize PtRu/C automatically. This method was proven to be a good fit for the mass production of PtRu/C. These benefits include effective, raw material saving, highly reproducible and labor-saving.

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INTRODUCTION Direct Methanol Fuel Cell (DMFC) is considered to be the most attractive power source for the transportation and portable electronic devices due to its high energy conversion efficiency [1-3]. However, the commercialization of DMFC is obstructed from the high cost of electrocatalysts. To reduce the cost of DMFC, the amount of noble metal catalysts should be reduced, while keeping the reactive site of metal catalysts remain unchanged or even higher. In order to obtain a maximum in catalytic performance while keeping a minimal mass of catalyst, noble metals are usually prepared in form of highly dispersed nanoparticles supported on conductive carbon supports [4-7]. So far, the impregnation, colloidal, and microemulsion methods are widely used for preparing the electrocatalysts for DMFC application [8-16]. Among these methods, the impregnation method seems to be an attractive way due to its simplicity. However, it also has problems with large scale applications because it usually includes complex and timeconsuming multi-step processes. Therefore, it is essential to improve the impregnation method to meet the need of the commercialization of DMFC. In addition, it is well known that pure platinum is readily poisoned by adsorbed carbon monoxide, an intermediate product of anodic methanol oxidation [17, 18]. At present, bimetallic PtRu catalyst is the preferred anode catalyst in the application of DMFC. Recent studies [19-25] revealed that there is numerous factors (e.g., morphology, structure, composition, uniformity, dispersion state, and alloying state) influence the physicochemical and electrocatalytic properties of the PtRu/C catalyst and there factors are usually affected by the preparation parameters. Here we present a systematic study of the influence of preparation parameters on the physicochemical and electrocatalytic properties of 40 wt% PtRu/C. X-ray diffraction (XRD), transmission electron microscope (TEM) and X-ray photoelectron spectroscopy (XPS) characterizations were carried out to determine the particle size and surface composition of the PtRu/C catalysts. Cyclic voltammetry (CV) experiments were also conducted to assess the electrocatalytic peroformance of the PtRu/C catalysts. The factors that affect the electrocatalytic performance of the PtRu/C catalysts toward methanol oxidation were analyzed. It is expected that these results will give a better understanding of the effects of preparation parameters on the electrocatalytic performance of PtRu/C catalysts for methanol oxidation. In addition, we introduced a new and highly reproducible method based on the flow injection (FI) technology for the preparation of the PtRu/C catalysts. This method was proven to be a suitable technology for the mass production of PtRu/C. These benefits include effective, raw material saving, highly reproducible and labor-saving.

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EXPERIMENTAL SECTION Chemicals All chemical reagents are analytical grade. H2PtCl6·6H2O, RuCl3, NaBH4, methanol, ethanol, formaldehyde, N2H4, isopropyl alcohol and KOH were obtained from Shanghai Chemical Regent Ltd. Vulcan XC-72 carbon was obtained from Cabot Company. All chemicals were used without further purification.

Preparation of PtRu/C Catalysts The carbon black, Vulcan XC-72 carbon from the Cabot Corporation was used as the support for the PtRu/C catalyst. Firstly, it was suspended in deionized water and sonicated. H2PtCl6 and RuCl3 were dissolved in deionized water and stirred. These two solutions were mixed together and stirred sufficiently. The pH of the mixed solution was adjusted by using 0.1 M NaOH or 0.1M HCl. Different reductant (NaBH4, N2H4, formaldehyde, methanol, ethanol and isopropyl alcohol) solutions of different concentrations were prepared in which the molar ratioes (α value) of reductant to Pt and Ru metals were 1, 5, 10, 20, and 50. The reductant solution was quickly poured into the pre-heated (95 oC) mixed solutions of the carbon and metal ions. After 20 min the prepared catalysts were filtered and washed with deionized water for several times. The catalysts were then dried in vacuum oven at 90 oC for 5 h. The metal loadings of catalysts were fixed at 40 wt % and the molar ratio of Pt to Ru is 1:1.

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Physical Characterization X-ray photoelectron spectroscopy (XPS) was measured with a Quantum 2000 Scanning ESCA Microprob equipment (Physical Electronics) using Al-Kα radiation. The binding energy was calibrated using a C 1s photoelectron peak at 284.6 eV as a reference. The surface composition was determined from the peak areas and the sensitive factors presented by Physical Electronics. The PtRu/C catalysts were characterized by powder x-ray diffraction (XRD) on an X’Pert PRO X-ray diffractometer with Cu-Kα radiation at 40 kV and 30 mA at room temperature. The 2θ angular regions between 20 and 90 o were recorded at a scan rate of 0.0167 0 for 10 seconds. Raman spectra were obtained using the Renishaw micro-Raman, model 1000, with Ar+ laser (514.5 nm). A Tecnai F30 transmission electron microscope (TEM) was employed for the analysis of the morphology, microstructure and composition of the PtRu/C catalysts.

Electrochemical Measurement Electrochemical experiments were carried out with a CHI 660A electrochemical workstation. A conventional, three-electrode cell consisting of glassy carbon (GC) with an area 0.1256 cm2 as the working electrode, Pt foil as the counter electrode and saturated

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calomel electrode (SCE) as the reference electrode, was used. The working electrode was prepared as follows. The GC electrode was polished carefully by using 0.3 and 0.05 μm alumina pastes. To prepare the catalyst ink, 14.4 mg of PtRu/C was ultrasonically dispersed in a mixture of 5 ml of ethanol and 1 ml of 0.1% Nafion solution for 30 min. 10 μl of the catalyst ink was then pipetted out on the top of the GC electrode and dried at 343 K, after that 10 μl of 0.1% Nafion solution was coated on it as a binder and dried at 343 K to yield a metal loading of 38 μg cm-2. For the electrochemical experiments, all the potentials in this paper were quoted against SCE and all solution were prepared by using ultra pure water (Millipore, 18 MΩ). The cyclic voltammetry (CV) curves were obtained in a voltage range from -0.241 to 0.56 V in 0.5 M H2SO4 with or without 1 M CH3OH at a scan rate of 10 mV s-1. The chronoamperometry tests were performed at 0.4 V in 0.5 M H2SO4 containing 1 M CH3OH solution. The CO stripping voltammograms were conducted with the three-electrode cell in 0.5 M H2SO4 solution. Firstly, nitrogen was purged to 0.5 M H2SO4 solution for 30 min, and then the adsorption of CO was performed by purging CO gas to 0.5 M H2SO4 solution for 30 min while maintaining the electrode potential at -0.18 V. After that, the nitrogen was purged again to the solution for 30 min to remove CO in the solution. Finally the CO stripping voltammograms were recorded between -0.241 and 0.56 V at a scan rate of 10 mV s-1 under N2.

RESULTS AND DISCUSSIONS

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Effect of Preparation Parameters on the Structure, Morphology and Surface Composition of Ptru/C Catalysts Fig. 1 and 2 show the X-ray diffraction patterns of the catalysts prepared using methanol as reductant. As can be found in Fig. 1 and 2, the peaks observed indicated the present of the face centered cubic (fcc) structure of Pt, represented by the crystalline planes (111), (200), and (220) (JCPDS, card 4-802). In addition, the 2θ values of the peaks for PtRu/C are shifted to higher values. This indicates that the PtRu solid solutions are formed. The absence of diffraction peaks of metallic Ru or RuO2 in the XRD patterns suggests that the Ru atoms are in the form of alloying with Pt, decorated on Pt particles or Ru being predominantly in a form of amorphous Ru-oxide [25]. Furthermore, only in the case of PtRu/C prepared at a pH of 14 (Fig. 2), the diffraction peaks are observed around 42o and 54o, those are attributed to the hexagonal structure of graphite (100) and (004). This result implies that the complete reduction of H2PtCl6 and RuCl3 is not achieved. Moreover, we calculate the lattice parameter (afcc) values and the Ru atomic fraction (XRu) in the PtRu alloy by using the following equations with deconvoluting the (220) diffraction peaks [26, 27].

α fcc =

2 × λ Kα 1 sin θ max

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(1)

Physicochemical and Electrocatalytic Properties of Ptru/C…

X Ru =

0.39155 − α fcc

297

(2)

K

where λ is the wavelength of X-ray (1.54056 Å), θ is the angle at the maximum of the diffraction peak (220). 0.39155 nm is the lattice parameter for pure platinum supported on the carbon and K=0.0124 nm is a constant. Table 1 shows the lattice parameter values and the corresponding variations in the Ru atomic fraction for the different PtRu/C catalysts. It can be obviously seen that the amount of Ru incorporation into the fcc structure of PtRu/C increase with the values of α and pH. Table 1. Characterization of the PtRu/C catalysts: Ru content (nominal, XRD) and average particle size (XRD) Catalyst

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PtRu/C (pH=10, α=1) PtRu/C (pH=10, α=5) PtRu/C (pH=10, α=10) PtRu/C (pH=10, α=20) PtRu/C (pH=10, α=50) PtRu/C (pH=4, α=50) PtRu/C (pH=6, α=50) PtRu/C (pH=8, α=50) PtRu/C (pH=14, α=50)

Ru content (Nominal, at.%) 50 50 50 50 50 50 50 50 50

Ru in alloy (XRD, at. %) 1.9 7.8 19.2 27.9 36.3 2.8% 17.2% 25.5% -

Average particle size (nm) 8.2 7.8 5.3 3.5 2.3 10.4 8.4 3.8 -

Figure 1. Patterns of X-ray diffraction of the PtRu/C prepared at different α values (methanol as reductant, pH=10)

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Figure 2. Patterns of X-ray diffraction of the PtRu/C prepared at different pH values (methanol as reductant, α= 50)

Figure 3. TEM images of the PtRu/C catalysts prepared at different pH values using methanol as reductant (a: pH=14; b: pH=10; c: pH=8; d: pH=7; e: pH=6; f: pH=4)

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Fig. 3 presents the typical TEM images of PtRu/C prepared at different pH values using methanol as reductant. For the catalyst prepared at the pH of 14 (Fig 3a), no particles is visible on the carbon support. It can be also observed in Fig 3b that well dispersed PtRu nanoparticles are formed on the carbon support. The mean size of PtRu particles is estimated to be 2.3 nm. In the case of the PtRu/C catalysts prepared at a lower pH value, larger particles (for PtRu/C (pH=4, α=50), PtRu/C (pH=6, α=50), PtRu/C (pH=7, α=50), and PtRu/C (pH=8, α=50) catalysts, the particle sizes are 12.1, 7.9, 4.1 nm, respectively) with some agglomerates of PtRu particles are formed on some areas of carbon (Fig 3c~f). The average particle sizes calculated from TEM are found to be consistent with those from the XRD peak widths. Moreover, the effect of α value on the morphology of PtRu/C catalysts is also investigated, and the results show that the particle size of the PtRu/C catalysts decreases with the increase of α value. A detailed discussion of the above-mentioned phenomena is presented in the following section. Table 2. XPS results of the as-prepared catalysts Catalyst

Species

PtRu/C (pH=10, α=50)

Pt 4f7/2

Ru3p3/2

PtRu/C (pH=14, α=50)

Pt 4f7/2

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Ru3p3/2

PtRu/C (pH=4, α=50)

Pt 4f7/2

Ru3p3/2

PtRu/C (pH=10, α=1)

Pt 4f7/2

Ru3p3/2

PtRu/C (pH=10, α=10)

Pt 4f7/2

Ru3p3/2

Pt0 Pt2+ Pt4+ Ru metal RuO2 RuO2·xH2O Pt0 Pt2+ Pt4+ Ru metal RuO2 RuO2·xH2O Pt0 Pt2+ Pt4+ Ru metal RuO2 RuO2·xH2O Pt0 Pt2+ Pt4+ Ru metal RuO2 RuO2·xH2O Pt metal PtO PtO2 Ru metal RuO2 RuO2·xH2O

Binding energy (eV) 71.3 72.5 74.9 461.8 463.4 466.2 71.4 72.5 74.7 461.8 463.6 465.7 71.3 72.5 74.8 461.8 463.5 466.0 71.3 72.5 74.7 461.7 463.5 465.9 71.2 72.3 74.7 461.8 463.5 465.7

Relative intensity 60.1% 23.5% 18.3% 55.3% 31.8% 12.9% 1.5% 7.0% 81.5% 6.1% 50.0% 43.9% 63.4% 20.6% 16.0% 67.2% 18.6% 14.2% 35.1% 26.1% 38.7% 25.6% 42.8% 31.6% 58.8% 27.6% 13.5% 60.0% 22.0% 18.0%

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Pt:Ru (at%)

51:49

35:65

53:47

29:71

54:46

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The surface compositions and chemical oxidation states of Pt and Ru in the PtRu/C catalysts are determined by XPS analysis. The amounts of Pt and Ru species are calculated from the Pt 4f and Ru 3p3/2 spectrum, respectively. These results are listed in Table 2. As can be seen from Table 2, the PtRu/C (pH=14, α=50) and PtRu/C (pH=10, α=1) catalyst show a prevailing presence of oxidation state species, whereas, the other three catalysts contain a larger fraction of metallic component. This indicates that the reduction of Pt and Ru ions can not be completely achieved at a higher pH and a lower α value.

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Mechanisms for Catalyst Impregnation

Figure 4. Effect of pH value on the particle size of the PtRu/C catalysts

The effect of pH value on the particle size of the PtRu/C catalysts is summarized in Fig. 4. The α value of the reductant in this study is fixed at 50. As can be seen in Fig. 4, the particle size of the PtRu/C catalyst reduced by NaBH4 increases with the increase of the pH value and approaches to a constant value when the pH value is above 8. However, the other catalysts exhibit an inverse trend: the particle sizes of the PtRu/C catalysts decrease with the increase of the pH value. This phenomenon could be explained from the influence of reduction rate on the nucleation. It is known that the formation of metallic nanoparticles on carbon follows a two-step scheme [28]. The reduction of metallic ions occurs earlier to form nuclei, and then nucleusmediated autocatalytic reduction of excess metallic ions is induced by reductant on the surface of nucleus. In the case of using NaBH4 as reductant, the reduction of PtCl62- is supposedly through the following reaction:

PtCl 62− + 2 BH 4− + 6 H 2 O → Pt + Cl − + 10 H 2 + 2 B(OH ) 3 [29, 30]

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At a high pH value, the reduction rate of PtCl62- is slow and only few nuclei are formed at the early period of the reduction. The reduction of PtCl62- at the latter period mainly occurs at the surface of the nuclei already formed instead of the formation of new nuclei and therefore led to the formation of larger particles. With the decrease of the pH value, the reduction rate is enhanced, and this favors the formation of much more nuclei and thus smaller average particle size. In the case of using isopropyl alcohol, ethanol, N2H4, formaldehyde and methanol as reductant, the reduction of PtCl62- via reactions (4) ~ (6):

PtCl 62− + NH 2 NH 2 → Pt + N 2 + 6Cl − + 4 H + [31]

(4)

PtCl 62− + H 2 O + HCHO → Pt + CO2 + 6Cl − + 4 H + [32]

(5)

3PtCl62− + 2 H 2O + nROH → 3Pt + mCO2 + 18Cl − + 12 H +

(6)

With the increase of pH value, the reduction rate is enhanced, and this leads to the formation of much more nuclei and thus smaller PtRu nanoparticles. The effect of α values and the kind of reductant on the particle size of PtRu/C catalysts is also investigated in this study. As indicated in Fig 5, it is found that the average particle size of PtRu/C catalysts decreases significantly with the increase of the α value, and maintain practically constant when the α value reaches 20. In addition, it can be seen that the reductants with higher reducing power, such as formaldehyde, NaBH4, methanol, N2H4, lead to smaller particles. These phenomena could be also attributed to the influence of reduction rate on the nucleation. In the case of the α value is relative lower and the reductant with lower reducing power, the reduction rate of metallic ions is slow and results in less nuclei at the very beginning of the reduction. Therefore, this leads to larger PtRu nanoparticles. The above results demonstrate that the pH value, the kind and amount of reductant has a significant influence not only on the particle size, but also on the surface composition of PtRu/C.

Figure 5. Effect of α value and kind of reductant on the average particle size of PtRu/C

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Electrochemical Measurements Fig. 6 displays the typical CO stripping voltammograms recorded at 10mV s-1 for the PtRu/C catalysts prepared by using methanol as reductant. In this work, CO stripping voltammograms are recorded twice for all the catalysts to obtain reliable and reproducible results. It can be observed in Fig. 6 that hydrogen desorption peaks are suppressed completely in the typical range of -0.24 to 0.1 V for all the catalysts, indicating that the Pt surface sites of these catalysts are saturated with COad species.

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Figure 6. CO stripping voltammograms of the PtRu/C catalysts recorded at 10 mV/s in 1.0 M H2SO4.

It is well known that the shape and the position of the CO oxidation peak are very sensitive to the composition of the PtRu/C catalysts. As can be seen in Fig. 6, the CO oxidation peak for the PtRu/C (pH=14, α=50) and PtRu/C (pH=10, α=1) catalyst is broader than those of the other four catalysts. This indicates a slow CO oxidation reaction. This may attributed to not only the lower amount of Pt and Ru loaded onto the carbon support but also the lower degree of PtRu alloy [33]. Table 3 gives the summary of the parameters such as the onset potential (Eon), peak potential (Epa), and electrochemical surface area (ECSA) extracted from the CO stripping voltammograms. As can be found in Table 3, the similarities of Epa and Eon values for PtRu/C (pH=14, α=50) and PtRu/C (pH=10, α=1) suggest that these two catalysts have a similar surface composition. The lowest Eon and Epa values for the CO oxidation reaction are observed for the PtRu/ C (pH=10, α=50). These results imply that this particular catalyst is catalytically more efficient than the other catalysts. This significant difference in the CO oxidation kinetics is likely due to the suitable amount of Ru incorporated into the fcc structure Pt and the suitable particle size of this catalyst. Furthermore, we measure the ECSA values for all the catalysts by calculating the COad oxidation charge after subtracting the back ground current with the assumption that 420 μC cm-2 as the oxidation charge of monolayer CO on a smooth Pt surface. The measured ECSA values are listed in Table 3. From Table 3, one can observe that the PtRu/C (pH=10, α=50) catalyst has the highest ECSA value. The high ECSA value means that a high number of

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exposed Pt atoms present on the electrode surface. The results may arise from the welldispersed PtRu particles on the carbon support. It also can be seen from Table 3 that the PtRu/C (pH=14, α=50) and PtRu/C (pH=10, α=1) catalyst has the smallest ECSA value. This may results from the incomplete reduction of Pt and Ru precursors and the blocking of surface metal sites by the large excess Ru oxides that are present in this sample according to the XPS results (Table 2.). Table 3.CO stripping characteristics for the PtRu/C catalysts

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Catalyst PtRu/C (pH=10, α=50) PtRu/C (pH=10, α=10) PtRu/C (pH=4, α=50) PtRu/ C (pH=14, α=50) PtRu/C (pH=10, α=1)

ECSA (m2 g-1) 92.6 79.2 70.8 41.3 45.2

Eon(V) vs. SCE 0.13 0.16 0.16 0.18 0.17

Epa(V) vs. SCE 0.26 0.30 0.35 0.39 0.39

Figure 7. Effects of pH (a) and α value (b) on the current density of the PtRu/C catalysts for the methanol oxidation reaction at 0.4 V

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To evaluate the electrocatalytic performace of the PtRu/C catalysts prepared using methanol as reductant, cyclic voltametric measurements of the methanol oxidation reaction are carried out in nitrogen saturated 0.5 M H2SO4 solution with 1 M CH3OH in the potential range from -0.241 to 0.56 V. The surface area normalized current density of methanol oxidation reaction at 0.4 V for the PtRu/C catalysts prepared at different pH and α values are shown in Fig 7. The results reveal that the PtRu/C prepared at higher pH and α values has better electrocatalytic performance for methanol oxidation, but it becomes less pronounced when the pH and αvalue is higher than 8 and 20, respectively. These differences can be attributed to the size effect and the different degree of PtRu alloying (Table 1).

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Application of Flow Injection Technique in Continuous Electrocatalyst Synthesis The impregnation reduction is widely used in the synthesis of catalyst due to its simpility. However, it also has problems with scale-up because it usually includes complex and timeconsuming multi-step processes. Flow injection (FI) technique is a simple, rapid, and versatile technique that receives widespread attention in the field of quantitative chemical analysis. In a typical flow injection system, the sample is injected into a continuous flow of reagent solution, dispersed, and transported to detector or reactor [34]. Flow injection analysis equipments are designed to deliver and react sample and reagents in the required order and ratios. Therefore, this technology is labor-saving and highly reproducible since there is no man-made interference during the whole process. A schematic diagram of the FI system for synthesis of the PtRu/C catalyst is shown in Fig.8. In the basic setup, the mixed solution of H2PtCl6, RuCl3 and carbon black is used as the carrier solution. Before the pH value of the carrier solution is adjusted by mixing with NaOH or HCl solution, it is preheated to 95 oC in an oil bath. Subsequently, the solution is mixed with the preheated reductant in the mixing zone. To ensure the complete reduction of metallic ions, the resulted catalyst is kept in the oil bath for more than 20 min. The PtRu/C catalyst is then filtered, extensively washed with deionized water and dried in vacuum oven at 90 oC for 5 h. In this setup, the pH and α values can be precisely controlled by adjusting the flow rates of the carrier, NaOH or HCl and reductant solution. Table 4 compares the electrocatalytic properties of two PtRu/C catalysts prepared by the FI system with the same operational parameters. As can be seen in Table 5, these two catalysts show identical electrocatalytic properties. This indicates that the FI technology is suitable for mass production of PtRu/C catalysts.

Figure 8. Schematic diagram of the FI system for the synthesis of PtRu/C catalysts

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Table 4. Comparison of the electrocatalytic properties of two PtRu/C catalysts prepared by the FI system with the same operational parameters Catalyst

ECSA (m2 g-1)

Eon(V) vs. SCE

PtRu/C-1 PtRu/C-2

95.2 94.1

0.14 0.14

Surface area specific activity at 0.4 V (mA cm-2) 0.21 0.22

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CONCLUSION In the present study, the influence of preparation parameters on the physicochemical and electrocatalytic properties of 40 wt% PtRu/C is investigated. The TEM and XRD results indicated that well-dispersed PtRu nanoparticles of around 2.3 nm in size can be obtained on the carbon support when a proper preparation condition is selected. The XRD and TEM measurements also reveal that the preparation conditions have a significant influence on not only the dispersion state, but also the amount of Ru incorporated into the fcc structure of PtRu/C. In addition, it is found that the reductants with higher reducing power, such as formaldehyde, KBH4, methanol, N2H4, lead to smaller particles and higher dispersion. On the other hand, the reductants with lower reducing power, such as ethanol, isopropyl alcohol, result in bigger particles and poorer dispersion. Moreover, the effect of pH value during impregnation was also investigated. In the case of using NaBH4 as a reductant, the favorable pH value is below 4, whereas in the case of methanol, N2H4 and formaldehyde as a reductant, the favorable pH value is higher than 8. The electrocatalytic activity of as-prepared catalyst was tested by cyclic voltammograms both in 1 M H2SO4 and 0.5 M H2SO4+ 1 M CH3OH. The results indicate that a reductant with higher reducing power, a high stoichiometric ratio of the reductant and a proper pH value are favorable to synthesize a catalyst with a higher degree of PtRu alloying, and thus a higher electrocatalytic activity toward methanol oxidation. In addition, a new and highly reproducible method based on the flow injection (FI) technology is also developed to automatically synthesize the PtRu/C catalysts. The results demonstrate that this method is suitable for the mass production of PtRu/C. These benefits include effective reaction, raw material saving, reproducibility and labor-saving.

ACKNOWLEDGMENTS This work was supported by the NSF of China (20673089), the 973 Program of China (2009CB939804), the Key scientific project of Fujian Province of China (2005HZ01-3) and the China Postdoctoral Science Foundation (No. 20070420120).

REFERENCES [1] Wilson M.S.; Gottesfeld S. J. Appl. Electrochem. 1992, 22, 1-7. [2] Surampudi S.; Narayanan S.R.; Vamos E.; Frank H.; Halpert G.; Laconti A.; Kosek J., Surya Prakash G.K.; Olah G.A. J. Power Sources 1994, 47, 377-385.

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[3] Ren X.; Wilson M.S.; Gottesfeld S. J. Electrochem. Soc. 1996, 143, L12-L5. [4] Park K.W.; Sung Y.E. J. Phys. Chem. B 2004, 108, 939-944. [5] Joo S.H.; Choi S.J.; Oh I.; Kwak J.; Liu Z.; Terasaki O.; Ryoo R. Nature 2001, 412, 169172. [6] Steigerwalt E.S.; Deluga G.A.; Cliffel D.E.; Lukehart C.M. J. Phy. Chem. B 2001, 105, 8097-8101. [7] Hyeon T.; Han S.; Sung Y.E.; Park K.W.; Kim Y.W. Angew. Chem. Int. Ed. 2003, 42, 4352-4356. [8] Zhou Z.H.; Zhou W.J.; Wang S.; Wang G.X.; Jiang L.H.; Li H.Q.; Sun G.Q.; Xin Q. Catal. Today 2004, 93-95, 523-528. [9] Watanabe M.; Uchida M.; Motoo S. J. Electroanal. Chem. 1987, 229, 395-406. [10] Han K.I.; Lee J.S.; Kim H.S. Electrochem. Acta 2006, 52, 1697-1702. [11] Shimazaki Y.; Kobayashi Y.; Yamada S.; Miwa T.; Konno M. J. Colloid Interf. Sci. 2005, 292, 122-126. [12] Solla-Gullon J.; Vidal-Iglesias F.J.; Montiel V.; Aldaz A. Electrochem. Acta 2004, 49, 5079-5088.. [13] Zhang X.; Chan K.Y. Chem. Mater. 2003, 15, 451-459. [14] Roucoux A.; Schulz J.; Patin H. Chem. Rev. 2002, 102, 3757-3778. [15] Liu H.; Song C.; Zhang L.; Zhang J.; Wang H.; Wikinson D. P. J. Power Sources 2006, 155, 95-110. [16] Ren L.; Xing Y. Electrochem. Acta 2008, 53, 5563-5568. [17] Burstein G.T.; Barnett C.J; Kucernak A.R.; Williams K.R. Catal. Today 1997, 38, 425437. [18] Hamnet A. Catal. Today 1997, 38, 445-457. [19] Dinh H.N.; Ren X.; Garzon F.H.; Zelenay P.; Gottesfeld S. J. Electroanal. Chem. 2000, 491, 222-233. [20] Vidakovic T.; Christov M.; Sundmacher K.; Nagabhushana K.S.; Fei W.; Kinge S.; Bönnemann H. Electrochem. Acta 2007, 52, 2277-2284. [21] Bock C.; Blakely M.A.; MacDougall B. Electrochem. Acta 2005, 50, 2401-2414. [22] Gavrilov A.N.; Savinova E.R.; Simonov P.A.; Zaikovskii V.I.; Cherepanova S.V.; Tsirlina G.A.; Parmon V.N. Phys. Chem. Chem. Phys. 2007, 9, 5476-5489. [23] Rojas S.; García-García F.J.; Järas S.; Martínez-Huerta M.V.; Fierro J. L. G.; Boutonnet M. Appl. Cata. A 2005, 285, 24-35. [24] Takasu Y.; Fujiwara T.; Murakami Y.; Sasaki K.; Oguri M.; Asaki T.; Sugimoto W. J. Electrochem. Soc. 2000, 147, 4421-4427. [25] Guo J.W.; Zhao T.S.; Prabhuram J.; Chen R.; Wong C.W. Electrochem. Acta 2005, 51, 754-763. [26] Liu Z.L.; Lee J.Y.; Chen W.X.; Han M.; Gan L.M. Langmuir 2004, 20, 181-187. [27] Yang B.; Lu Q.; Wang Y.; Zhuang L.; Lu J.; Liu P.; Wang J.; Wang R. Chem. Mater. 2003, 15, 3552-3557. [28] Chen D.H.; Wu S.H. Chem. Mater. 2000, 12, 1354-1360. [29] Glavee G.N.; Klabunde K.J.; Sorensen C.M.; Hadjipanayis G.C. Langmuir 1993, 9, 162169. [30] Glavee G.N.; Klabunde K.J.; Sorensen C.M.; Hadjipanayis G.C. Langmuir 1992, 8, 771773. [31] Deivaraj T.C.; Chen W.X.; Lee J.Y. J. Mater. Chem. 2003, 13, 2555-2560.

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[32] Li W.Z.; Liang C.H.; Zhou W.J.; Qiu J.H.; Zhou Z.H.; Sun G.Q.; Xin Q. J. Phys. Chem. B 2003, 107, 6292-6299. [33] Gasteiger H.A.; Markovic N.; Ross P.N.; Cairns E.J. J. Phys. Chem. 1994, 98, 617-625. [34] Ruzicka J.; Hansen E.H. Flow Injection Analysis 2nd ed.; John Wiley & Sons: New York, 1988: pp 15-31.

REVIEWED BY

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Prof. Zhong-Qun Tian at State Key Lab of Phys. Chem. of Solid Surfaces, Xiamen University (email: [email protected]).

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In: Polymer Electrolyte Membrane Fuel Cells… Editors: R. Esposito

ISBN: 978-1-60692-773-1 ©2009 Nova Science Publishers, Inc.

Chapter 11

CONDUCTING POLYMERS USED AS CATALYST SUPPORT FOR FUEL CELL APPLICATION Lei Li∗ and Jun Yang∗∗ School of Chemistry and Chemical Engineering, Shanghai Jiao Tong University Shanghai 200240, P. R. China

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ABSTRACT An ideal catalyst support for fuel cell application has not only high electron and proton conductivity, but also good permeability to gases and water. However, carbon, the most common catalyst support used in fuel cell, is impermeable to gases (oxygen, hydrogen and water vapor) and non-conductive to protons, which results in low catalyst utilization and poor catalyst performance. Recently, more attention has focused on the use of conducting polymers (e.g. polyaniline, polypyrrole and polythiophene) as catalyst supports and the promising results have been obtained. Combined with our recent research results, a review on conducting polymers and conducting polymers-carbon composites used as catalyst support for fuel cell application will be presented in this chapter.

INTRODUCTION Fuel cells convert the chemical energy of a fuel directly into electricity. Because fuel cells operate without a thermal cycle, they offer a quantum jump in the energy-conversion efficiency. Fuel cells also require no emission-control devices that are necessary in conventional energy-conversion devices. In order to meet global energy demands after fossil fuel depletion, and to control environmental pollution, fuel cells are attracting increasing attention in recent decades due to the high efficiency and low/zero emission. The six generic fuel cells in various stages of development are (1) phosphoric acid fuel cells (PAFCs), (2) alkaline fuel cells (AFCs), (3) polymer electrolyte membrane fuel cells or proton exchange ∗

[email protected]; ∗∗ [email protected]

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membrane fuel cells (PEMFCs), (4) molten carbonate fuel cells (MCFCs), (5) solid oxide fuel cells (SOFCs), and (6) direct methanol fuel cells (DMFCs). Among them, PEMFCs and DMFCs are being investigated and developed towards various applications for low/zeroemission vehicles, portable electronic devices, and mobile and stationary power generation [1]. The structure of PEMFCs and DMFCs is quite similar, as shown schematically in Figure 1. The electrochemical reaction takes place in the membrane electrode assembly of the fuel cell. Each membrane electrode assembly consists of electrodes (anode and cathode) with a thin layer of catalyst, bonded to either side of a proton exchange membrane (PEM). Fuels including hydrogen for PEMFC or methanol aqueous solution for DMFC flow through channels in the flow field plate to the anode where catalyst promotes them separation into protons, electrons or CO2 produced in the case of DMFC. Air flows through channels in the flow field plate to the cathode. Oxygen in the air attracts the protons through the PEM, while electrons are forced through an external circuit. Combining the anode and cathode reactions, the overall cell reaction is: PEMFC system: H 2 + 1 O2 → H 2O

2

DMFC system: CH 3OH + 3 O2 + H 2O → CO2 + 3H 2O

2

O2

CH3OH + H2O

H2O

Diffusion Layer

Catalyst Anode

Membrane

Catalyst Cathode

+

Flow Field Plate

Diffusion Layer

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H+ / H2O

Flow Field Plate

-

CO2

Figure 1. Schematic diagram of a direct methanol fuel cell.

In the path towards fuel cell commercialization, however, one of the big challenges is to reduce the material cost. One of the major contributors to the high cost is the platinum based catalyst. In order to reduce the cost of the fuel cell catalysts, one approach is to explore new catalyst materials. Various Pt-M (M=Ru, Co, Cr, Fe, Ni, Cu…) alloys or non-noble catalysts are being intensively investigated to increase the catalytic activity [2-5]. Another approach is to develop new catalyst support materials to achieve high dispersion, utilization, activity and stability for catalysts [3]. This is particularly important for lowering the fuel cell cost by

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Conducting Polymers Used As Catalyst Support For Fuel Cell Application

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reducing the amount of expensive Pt-based catalysts or enhancing the reactivation of nonnoble catalysts. Carbon is the most common catalyst support for PEMFC and DMFC systems. Nowadays the typical design of catalysts is to disperse and fix the nanosized noble metal (or multimetallic mixture/alloy) particles on the carbon support. The primary role of the carbon support is to provide electrical connection between the widely dispersed catalyst particles and the porous current collector (carbon cloth or paper). However, because it is impermeable to gases and does not conduct protons, the incorporation of carbon into the catalyst layer restricts gas (hydrogen and oxygen), water and proton transport, and limits the achievable performance of the catalyst and results in low catalyst utilization and high material cost. An ideal catalyst support should be gas and water permeable, and conduct both protons and electrons. Such a material could replace both carbon and polyelectrolyte binder (Nafion® perfluorinated ionomer) in the catalyst layer of the membrane electrode assembly, and improve the cell performance. Recently, numerous attempts have been made to utilize conducting polymers (CPs), such as polyaniline (PAn), polypyrrole (PPy), polythiophene (PTh) and poly(3,4dioxyethylenethiophene) (PEDOT) shown in Figure 2, as catalyst supports for hydrogen and methanol fuel cell applications. Conducting polymers are porous with high specific surface area and permeable to gases and water, and they exhibit both ionic and electronic conductivity. Then, only a two-phase boundary is needed for electron and proton transfer for the electrochemical reaction, compared to the three-phase boundary when carbon is used as the catalyst support. Moreover, the nature of CPs will retard CO poisoning of the noble metal catalysts. For these reasons, utilization of the noble metal catalysts can be enhanced. In addition, some conducting polymers, such as polypyrrole, can be used as a matrix to entrap some cheap transition metal atoms and generate active sites for oxygen reduction. It also provides the possibility of making a variety of non-noble metal catalysts towards the oxygen reduction. Therefore, conducting polymers as a promising catalyst support have attracted more research interests during the last decade. H

[

H

H

H

N

N

N

]n

N

[

Polyaniline

[ S S

Polythiophene

N

H

H

Polypyrrole O

[

N

S

]n

O

O

O

S S

S O

]n

O

Poly(3,4-dioxyethylenethiophene)

Figure 2. Chemical structures of conducting polymers

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Lei Li and Jun Yang

CONDUCTING POLYMERS The structure and morphology of conducting polymer depends on the preparation method. Conducting polymers can be obtained either by electrochemical polymerization from an electrolyte solution or by homogeneous chemical reaction with an oxidizing agent (e.g. FeCl3). Whereas electrochemical polymerization leads to the formation of thin films with a thickness of few micrometers on an electrode surface, a fine-grained material is obtainable by chemical oxidation in bulk solution. For practical applications, small particles have the advantage of a large specific surface area improving the overall reaction rate, the degree of conversion, and the space–time yield.

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Electrochemical Preparations Electrochemical synthesis of conducting polymers is a well-established technique that can successfully incorporate a wide spectrum of electrolyte anions soluble in the polymerization medium, into the growing chain of conducting polymers. At the same time electrochemistry has provided equally efficient techniques for incorporating insoluble particles of metal or alloy, etc. into growing polymer films, giving birth to a number of hybrid nanocomposites. In the fuel cell, it requires both the catalyst particles and their supports to be highly stable, for building a successful electrode. Therefore, appropriate metal particles dispersed in conducting polymers present a promising catalytic material for fuel cell applications. Catalytic particles dispersed along the entire thickness of the electrode lead to poor utilization of the catalyst. Therefore, conducting polymer bed with catalyst particles gathering near the surface presents an optimum situation in this respect. In general, conducting polymer film will be electrochemically polymerized onto the substrates (e.g. Pt or Au, glassy carbon, stainless steel) from an electrolyte solution, and then nanosized metal catalysts will be electrochemically deposited on the conducting polymer films. This was accomplished by using constant potential, galvanostatic or pulse galvanostatic, cyclic voltammetric, or a repetition square wave potential methods. This principle has been followed in a good number of works describing the preparation of nanocomposites, in which nanoparticles of Pt or Pt-M (M=Ru, Sn, Os, Mo) alloy are combined with conducting polymers, such as PAn, PPy and PTh (see Table 1). PAn and PPy are the most common conducting polymers used as support for the metal nanoparticles forming a composite catalyst for oxidation of methanol [6-10, 14-18, 20,21, 2327] and hydrogen [11, 14], and reduction of oxygen [22]. They have been mostly synthesized by a galvanostatic method (GM), a cyclic voltammetric method (CVM) or a potentostatic method (PM) [6-8, 10-22]. Polyaniline and PPy prepared by the above methods had a granular structure. Compared with CVM and PM methods, the pulse galvanostatic method (PGM) is a modified galvanostatic technology which may supply instantaneously high overpotential regardless of the relatively small mean current. The discontinuous current process reduces the concentration polarization of electrode/solution interface. With the PGM technique, PAn nanofibers would be prepared [9]. Compared with granular PAn film, PAn nanofibers film has better conductivity and higher specific surface area. The larger genuine area in the same geometric area is beneficial for Pt dispersion. Furthermore, the smaller

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transfer resistance of conductive particles in fibular PAn film decreases the catalyst poison [28]. The PAn nanofibers electrode modified by Pt microparticles, at the same Pt loading, exhibits a considerably higher catalytic activity towards the methanol oxidation than that of the granular PAn electrode modified by Pt particles. Template-assisted electrochemical synthesis is an effective method to prepare conducting polymeric nanotubes, nanofibers or nanocones of polypyrrole [29], polyaniline [30], poly(ophenylenediamine) [31] and poly(3-methylthiophene) [32]. The detailed template-synthetic procedures of conducting polymer nanostructures and the composites deposited with metal nanoparticles on Nafion and graphite electrodes are shown schematically in Figure 3. Nafion solution, coated on the graphite electrode, not only acts as binder, but also provides both ionic and electronic contact and favors proton transport. The alumina membranes are hot-pressed with the graphite that will be used as a current collector during the electrochemical polymerization. Through controlling the synthesized conditions, polypyrrole, polyaniline, poly(o-phenylenediamine) nanotubes [29-31], or poly(3-methyl) thiophene nanocones [32] can be obtained. Typical polypyrrole nanotubes and poly(3-methyl) thiophene nanocones are shown in Figure 4 and Figure 5, respectively. It has been reported [29-32] that these conducting polymers have enhanced conducting and charge transport properties compared to the conventionally synthesized conducting polymers. Platinum deposited on these conducting polymers nanostructure exhibited better catalytic activity and stability than on the carbon black and on the conventional conducting polymers mentioned before (without template). Graphite electrode (Nafion coating)

Electrochemical

Conducting polymer nanostructure

Electrodeposited

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Polymerization

Alumina membrane

Metal nanoparticles

Template Dissolution

Catalyst-conducting polymer composite

Figure 3. Template-assisted electrochemical synthesis of nanostructure conducting polymer.

Figure 4. (a-b) SEM images of polypyrrole nanotubes [29].

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Table 1. Electrochemical preparation of catalyst-conducing polymers electrodes for fuel cell applications Substrate Pt or Au wires

Conducting polymer PAn

Electrochemical polymerization method Cyclic voltammetric method

Au wire

PAn

Cyclic voltammetric method

Stainless steel

PAn

Glassy carbon

PPy

Stainless steel Au electrode Glassy carbon Au sheet Au sheet Au sheet

PPy PAn PAn PPy and PTh PAn PAn

Glassy carbon Glassy carbon Pt or indium-tin oxide Au electrode

PAn PPy PAn

Galvanostatic or pulse galvanostatic methods Linear sweep voltammetric method Galvanostatic method Cyclic voltammetric method Cyclic voltammetric method Galvanostatic method Cyclic voltammetric method Cyclic voltammetric method Cyclic voltammetric method Potentostatic method Cyclic voltammetric method

PEDOT*

Potentostatic method

Pt sheet Glassy carbon

PAn PAn

Cyclic voltammetric method Potentostatic method

Catalyst

Electrodeposited method

Reference

Pt, Pt-Ru

Potentostatic or repetition square wave potential signal methods Repetition square wave potential signal method

6,7

9

Pt

Cyclic voltammetric galvanostatic methods Potentostatic method

Pt, Ru, Ir Pt, Pt-Ru, Pt-Sn Pt Pt Pt Pt-Ru

Galvanostatic method Potentostatic method Potentostatic method Pulse galvanostatic method Potentostatic method Potentostatic method

11 12 13 14 15 16

Pt-Sn Pt Pt

Cyclic voltammetric method Cyclic voltammetric method Cyclic voltammetric method

17 18 19

Cyclic voltammetric and potentostatic method** Potentostatic method Potentostatic method

20

Pt, Pt-Ru, Pt-Os, Pt-Mo, Pt-RuMo, Pt-Ru-Os Pt

Pt, Pt-Pb, Sn** Pt Pt

Pt-

or

pulse

8

10

21 22

PEDOT: poly(3,4-dioxyethylenethiophene), ** Pt deposited onto the electrode by firstly cyclic voltammetry and followed by cathodic polarization at the constant potential, and then Pb and Sn deposited onto the Pt by cathodic polarization at the constant potential.

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Figure 5. HR-TEM image of poly(3-methyl) thiophene nanocones [32].

A repetitive square wave potential signal (RSWPS) is conceived as a suitable method to obtain a better scattering of metallic particles on conducting polymer film electrodes [33], thus improving the catalytic properties of the modified electrodes for methanol oxidation. Applying the RSWPS technique, different metal catalysts, including Pt, Pt-Ru, Pt-Mo, Pt-Os, Pt-Ru-Mo, Pt-Ru-Os, were electrodeposited on PAn/Au or Pt films to obtain catalytic electrodes for methanol oxidation [6-8]. Hepel [34] proposed a new method for the formation of composite polypyrrole films containing a highly dispersed three-dimensional array of platinum catalyst particles. PtCl42anions, which were trapped in polypyrrole matrix during the electrochemical polymerization of pyrrole, were reduced to elemental platinum particles and incorporated in the electrically conducting polypyrrole films. This method permitted the synthesis of films with uniform distribution of small platinum particles with an average size of 10nm and showed better catalytic activity than two-dimensional films because of its high surface area. Trueba et al [11] electrosynthesized PPy films of various thicknesses containing particles of noble metals such as Pt, Ru and Ir on the surface of austenitic stainless steel substrate. The catalytic activity of these materials has been investigated for hydrogen oxidation. Recently, Kitani et al [22] found that the catalytic activity of Pt dispersed PAn electrode towards oxygen reduction was almost twice as high as that of Pt dispersed carbon electrode. Yano et al [35] investigated the electro-oxidation of methanol on platinum alloys with metal (Cr, Ni, In, Co, Sb, Bi, Pb, and Mn) dispersed on the PAn films by cathodic electrodepostion, and found the highest activity was obtained for the Pt-Sn particles on the PAn films. Some research groups [8, 12, 13, 36] also investigated the electro-oxidation of carbon monoxide on platinum-based alloys, including Pt, Pt-Ru and Pt-Sn, incorporated into PAn. Other kinds of conducting polymers, such as poly(3,4-dioxyethylenethiophene) [20], poly(ophenylenediamine) [31, 37] and poly(N-acetylaniline) [38] have been also electrochemically polymerized, and investigated as catalyst support for methanol oxidation.

Chemical Preparations Compared to the electrochemical polymerization of conducting polymers, there are several advantages about the chemical polymerization of conducting polymers as catalyst

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support for fuel cell applications. (1) the preparation is very simple, and suitable for the mass production of materials, (2) this processing can incorporate polyelectrolyte such as polyvinylsulfate and poly(styrene-4-sulphonate) into the conducting polymer matrix during the polymerization, which would promote ion conductivity and mechanical stability of conducting polymers. In general, conducting polymers are synthesized by in situ chemical polymerization with various kinds of oxidants (cationic oxidant: Fe3+, or anionic oxidant: S2O82-) containing polyelectrolyte solutions. The summary of conducting polymer-polyelectrolyte composites as catalyst support for fuel cell application is listed in Table 2, together with key properties and comments. Pickup et al [39-42, 46, 50, 51] have prepared lots of conducting polymerspolyelectrolyte composites and researched their characterizations as catalyst support for fuel cell application. All composites were prepared by chemical polymerization of monomer in the presence of the polyelectrolyte in water or aqueous acetonitrile, by using Fe3+ to oxidize the monomer. Synthesis conditions and compositions of these composites are given in Table 3. Different morphologies of these composites were obtained. PPy-PSS can be formed as spherical particles, whose size can be controlled from tens of nm to ca. 1µm simply by changing the reagent concentrations [50]. At high concentrations the small particles generated agglomerate into a sponge-like morphology. PEDOT-PSS [39] and PAn-PSS also have a sponge-like morphology. PPy-Nafion appears to be gel-like with no discrete particles. The catalytic activity of metal nanoparticles such as Pt [40-42], Pt-Ru [41] and PtO2 [51] dispersed on these composites was investigated. Oxygen reduction properties of Pt/PPy-PSS and PtO2/PPy-PSS nanocomposites were prevailing while electrical conductivity suffered remarkably. This difficulty was successfully overcome to a large extent by using PEDOT as the conducting polymer support [42]. EDOT was polymerized in the presence of NaPSS aqueous solution and the resultant composite (electrical conductivity 9.9S cm-1) was stirred in aqueous formaldehyde (18%) solution of H2PtCl6 to deposit Pt nanoparticles (~4nm) on the composite particles. This treatment resulted in a lowering of the conductivity of the catalyzed material to 4S cm-1. This catalyst works well as electrode of fuel cell and its performance is comparable to the commercial carbon supported catalyst at a higher Pt loadings. However, there is a potential problem on the catalysts dispersed on conducting polymer-polyelectrolyte composites prepared by Pickup team. Since the strong association of the Fe3+ with the polyelectrolyte, it caused most of the conducting polymer-polyelectrolyte composites to retain Fe3+ even after washing process. Most of these composites were found to contain Fe3+ associated with the polyelectrolyte [39]. These iron ions will be absorbed into the proton exchange membranes through ion exchange reaction during the operation of fuel cell, and thus decrease the proton conductivity of membranes. Finally, it will lower the performance of fuel cell, including the durability and stability. In order to solve this problem, the anionic oxidant, S2O82-, was used as oxidant in the polymerization of conducting polymers [45, 4749]. When Na2S2O8 is being used as the oxidant, a fine grained PPy-PSS nanocomposite with size of 20-45nm was obtained [49]. Such nanocomposite formed by mixing an aqueous solution of 0.1M pyrrole and 0.1M sodium polystyrenesulfonate (NaPSS) with an aqueous solution of 0.1M Na2S2O8. This nanocomposite has a homogeneous particle size distribution. When Pt nanoparticles were dispersed on this PPy-PSS support, the catalyst showed a good activity for methanol oxidation.

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Table 2. Summary of conducting polymer-polyelectrolyte composites as catalyst support for fuel cell application Conducting polymer

Oxidant

Polyelectrolyte

Catalyst

Comments

Reference

PEDOT

FeCl3, Fe(NO3)3, Na2S2O8

PSS

Pt, Pt-Ru

Pt catalysed samples favor for O2 reduction and methanol oxidation.

39-45

Pt-Ru catalysed samples are effective for methanol oxidation. PEDOT

Fe(NO3)3

PVS

Pt, Pt-Ru

Effective for O2 reduction and methanol oxidation.

41

PAn

FeCl3

PSS

Pt

Poor O2 reduction performance with Pt catalysed samples.

46

PPy

(NH4)2S2O8

DEHS

Pt-Ru

Good catalytic activity for methanol oxidation at room temperature

47

PAn

(NH4)2S2O8

DBSA

Pt-Ru

Effective for methanol oxidation.

48

PPy

FeCl3, Fe(NO3)3, Na2S2O8

PSS

Pt

Effective for methanol oxidation and oxygen reduction.

46, 49

PPy

FeCl3

Nafion

Soluble in ethanol; disperses in water

46

PEDOT

FeCl3

Nafion

Soluble in ethanol; disperses in water

46

afion: perfluorosulfonic acid, PSS: poly(styrene-4-sulphonate), PVS: polyvinylsulfate, DEHS: di(2-ethylhexyl) sulfosuccinate, DBSA: dodecylbenzene sulfonic acid

318

Lei Li and Jun Yang Table 3. Synthesis conditions, compositions and yields for conducting polymerpolyelectrolyte composites.

Composite

Solvent

Fe3+ to monomer ratio

Monomer: polyelectrolyte

Reactio n (hr)

% Yield (S cm-1)

Ref.

PPy-PSS

H2O

5

3.3

0.5

100

40, 50

PEDOT-PSS

CH3CN/H2O

4

7.5

70

93

39

PEDOT-PSS

H2O

5

5

2

62

39

PEDOT-PSS

H2Oa

5

5

2

100

51

PAn-PSS

H2O

2.5

6

4.5

-

46

PPy-Nafion

CH3CN/H2O

10

8

1

~100

46

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a. Low volume, emulsion polymerization; -: not reported

PAn has four types of redox states: the leucoemeraldine base (reduced form), the emeraldine base (EB, half-oxidised form), the emeraldine salt (ES, half-oxidised, protonated form) and the pemigraniline base (fully oxidized form). Utilizing the variation of these redox states in acidic and neutral media, Unwin et al [52] successfully prepared Pt/PAn composite by a simple two-step method. First PAn (EB) and K2PtCl4 were mixed in formic acid. The treatment of the EB form with acid produces the doped conducting ES form. Pt particles can be observed both at the surface of PAn film and embedded within it. Then, the Pt/PAn film formed was immersed into a neutral aqueous solution of K2PtCl4 for further Pt deposition. The second Pt deposition step appears to cause a leveling of the surface topography. In both the steps the driving force for Pt formation is that PAn (ES) acts as an electron donor and the platinum salt as an electron acceptor, resulting in the oxidation of PAn (ES) to the pernigraniline form and the reduction of K2PtCl4 to metallic Pt. In acidic media, the pernigraniline form may undergo reduction back to the ES form [53], thereby promoting the further reduction of the Pt complex to metallic Pt. Recently, Li [54] prepared polypyrrole nanofibers with a 20 nm diameter by an interfacial polymerization without the use of a template or any dopant. The monomer and the oxidant Na2S2O8 are initially separated by the boundary between the aqueous and organic phase. The polymerization occurs primarily at this interface, where all the components needed for polymerization come together (Figure 6 a, b). The pure polypyrrole nanofibers first forms at the interface. In a conventional synthesis, these nanofibers are exposed immediately to unreacted monomer and oxidant. In contrast, in an interfacial polymerization, these nanofibers rapidly move away from the interface and diffuse into the water layer, as shown in Figure 6 c, d. In this way, the nanofibers pull themselves away from the reaction front, thus avoiding overgrowth and allowing new nanofibers to grow at this interface. After washing and drying processes, the PPy nanofibers were obtained as shown in Figure 7 a. Pt particles with 3-4nm size were deposited on the prepared PPy nanofibers by chemical reduction

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method. The Pt/PPy nanofiber catalyst shows a higher electrochemical active surface area and higher methanol oxidation catalytic activity than the commercial Pt/C.

Figure 6. Snapshots showing the interfacial polymerization of pyrrole. The reaction times are (a) 0 sec, (b) 30 sec, (c) 90sec, (d) 1h, (f) PPy powders. The top layer is an oxidant aqueous solution and, the bottom layer contains pyrrole dissolved in an organic solvent.

-2

25

Current density / mA cm

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Figure 7. SEM images of PPy nanofibers and Pt/PPy nanofibers catalyst.

JM 40% Pt/C 40% Pt/PPy nanotubes

20 15 10 5 0 -5 -10

-200

0

200

400

600

800

1000

E / mV vs. Ag-AgCl Figure 8. Cyclic voltammograms of 40% Pt/PPy nanofibers and commercial 40% Pt/C. Electrolyte: 0.5M H2SO4 + 1.0M CH3OH. Scan rate: 50mV s-1, temperature: 25oC. Pt loading: 500μg cm-2.

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pyrrole

PPy

Using sprayed layer-by-layer technique, Michel et al [55] fabricated a high membrane electrode assembly (MEA) for PEMFC from polyaniline nanofibers prepared by an interfacial polymerization method. Pt nanoparticles were dispersed on the PAn nanofibers by chemical reduction method. The MEA was fabricated by sprayed layer-by-layer (LBL) assembly: Pt/PAn nanofiber catalyst dispersion in isopropanol was sprayed on the Nafion membrane used as the first LBL component. Nafion solution in isopropanol was used as second partner for LBL assembly. Then Pt/PAn nanofiber and Nafion solutions were alternately sprayed on the membrane, until the desirable thickness was obtained. Through charge transfer interaction between Nafion and Pt/PAn nanobfiber, an LBL assembly on Nafion membrane is formed. Microscopic investigation reveals that PAn nanofibers have a strong interdigitation and thus create a dense and extensive 3D network of fibers. This morphology increases the probability of contacts between the strands, while substantial porosity also offers efficient mass transport. Thus, the Pt/PAn nanofiber LBL MEA demonstrated a power density of 63mW cm-2 and yielded a Pt utilization of 437.5W g-1 Pt, which is comparable to the traditional fuel cell using carbon as Pt support. However, the amount of Pt used here is only half lower than for usual carbon-support Pt catalysts. Recently, a new MEA based on PPy modified Nafion membranes as catalyst support for DMFC application was proposed by Li [56]. Using a transport controlled chemical polymerization method described schematically in Figure 9, polypyrrole films were formed on the surface of Nafion membranes. The Nafion membrane separates the aqueous monomer solution from the solution containing oxidant. The monomer permeates the membrane from one side and the oxidant from the other side, and PPy forms where the reactants meet. From the micrographs in Figure 9, it can be seen that a thin PPy film is mainly formed at the surface of membrane, which is in contact with the oxidant. The thickness of PPy film can be easily controlled by the reaction time. Furthermore, Pt nanoparticles were deposited on the PPymodified Nafion membrane by direct chemical reduction method. SEM images in Figure 10 show that Pt nanoparticles uniformly exist in the PPy-Nafion membrane with a porous structure. The power density of 12.25mW cm-2 was obtained in DMFC with 0.67mg cm-2 and 0.56mg cm-2 of Pt on the anode and cathode, respectively.

Nafion

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320

Na2S2O8

δ Figure 9. Preparation of PPy modified Nafion membrane as catalyst support.

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Figure 10. SEM of Pt/PPy-Nafion membrane electrode.

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Besides the conventional conducting polymers mentioned above, some new kinds of conducting polymers such as poly(N-vinyl carbazole) and poly(9-(4-vinyl-phenyl) carbazole) (see Figure 11) have been synthesized as catalyst support for fuel cell applications [57].

Figure 11. Chemical structures of (a) poly(N-vinyl carbazole) and (b) poly(9-(4-vinyl-phenyl) carbazole)

Conducting Polymers-Carbon Black Composites Carbon black such as Vulcan XC-72 and Ketjen Black-EC 300J has been used as catalyst support broadly and commercially for fuel cell application. However, the rate of methanol oxidation is low partly due to low platinum utilization on this conventional carbon black support, which is, in turn, related to the low electrochemically accessible surface area for the deposition of Pt particles. Moreover, carbon black particles generally contain sulfur groups and may cause Pt particles aggregation [58]. Conducting polymers used as catalyst support would improve the utilization of metal catalysts due to their specific characterizations. However, their low conductivity compared to carbon and their poor stability during catalyst working are unfavorable for catalyst support. Considering the advantageous combination of

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carbon black and conducting polymers, more and more attention has been paid to conducting polymer-carbon composites. Vulcan XC-72 carbon black particles were incorporated into PAn matrix by an electrochemical codeposition technique during the electrochemical polymerization of aniline monomer [59]. The incorporation of carbon particles leads to the improvement of the physical and electrochemical properties of the PAn film. Likewise, the PAn film also provides a suitable matrix for conventional carbon black particles, and forms composite support with increased electrochemical accessible surface areas and decreased charge transfer resistance. Compared to PAn support, PAn-carbon black composite with dispersed Pt and Pt-Ru particles exhibited higher catalytic activity for methanol oxidation due to better nanosized catalyst dispersion and utilization. A similar behavior has also been obtained in the RuxMoySez/PAncarbon black catalysts for oxygen reduction reaction even in the presence of methanol [60]. Recently, we have prepared PPy-carbon black (Vulcan XC-72) composite by in situ chemical polymerization of pyrrole monomer on carbon particles using ammonium persulfate as oxidant [61]. Pt nanoparticles were deposited on the prepared polypyrrole-carbon composites by chemical reduction method. The influence of naphthalene sulfonic acid (NSA) as dopant on the characterization of support and the catalytic activity of Pt/PPy-XC72 catalysts was investigated in detail. From the TEM measurements shown in Figure 12, it can be observed that well dispersed, spherical platinum particles with size ranging from 3 to 4nm were anchored on the surface of PPy-XC72 composite supports. Calculated from Pt (220) peak in Figure 13 by means of the Scherrer formula, the average size of Pt particles on PPyXC72 with and without NSA as dopant was 3.3nm and 3.6nm, respectively.

Figure 12. TEM images of Pt/PPy-XC72 with (a) and without (b) NSA as dopant.

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Figure 13. X-ray diffraction of Pt/PPy-XC72 with (a) and without (b) naphthalene sulfonic acid as dopant.

The electronic conductivity of different catalyst supports was measured via four probe measurement method. The electronic conductivity of carbon black (Vulcan XC-72), PPyXC72 with and without dopant was 28S cm-1, 25.9S cm-1 and 12 S cm-1, respectively. A high electronic conductivity of PPy-XC72 with NSA as dopant should be due to that NSA, as a macromolecule acid, can be doped into the 3D structure of PPy matrix and form a proton acid structure, which improves the ability of electronic transfer. The electrochemical reactivity and electrochemical active surface areas of different catalysts were determined by cyclic voltammetry. As shown in Figure 14 (I), the hydrogen adsorption and desorption peaks of Pt/PPy-XC72 with NSA as dopant (a) are bigger than those of the other two electrochemical catalysts. The electrochemically active surface of 40% Pt/PPy-XC72 (85m2 g-1Pt) with NSA as dopant is higher than that without NSA (68m2 g-1Pt) and commercial 40% Pt/C (62m2 g-1Pt). The catalytic activity towards methanol oxidation of different catalysts is shown in Figure 14 (II). It can be found that the onset potential for methanol oxidation of Pt/PPy-XC72 with dopant has a negative shift with a higher peak current density compared to others. It indicates that the excellent methanol oxidation ability is obtained for the Pt/PPy-XC72 catalyst with NSA as dopant. In addition, the CO stripping experiments reveal that Pt/PPy-XC72 catalysts have good CO-tolerant ability due to incorporation PPy into the catalyst support. High cost of noble metal catalysts is one of the major challenges for the practical development of PEMFC and DMFC. Some strategies have been taken to overcome this limitation. First is to develop new catalysts by using a proportion of transition metals such as Fe, Co to replace noble metals. Second is to develop low-cost non-platinum metal catalysts with similar activity and durability to the currently used Pt-based catalysts.

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(I)

(II)

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Figure 14. Cyclic voltammograms of 40wt% Pt/PPy-C with (a) and without (b) naphthalene sulfonic acid as dopant and commercial 40wt% Pt/C (c). Electrolyte: (I) 0.5M H2SO4 and (II) 0.5M H2SO4 + 1.0M CH3OH. Scan rate: 20mV s-1, temperature: 25oC, and the Pt loading: 200µg cm-2.

Figure 15. X-ray diffraction of (a) commercial 40% Pt/C and (b) 25%Pt-10%Fe/PPy-XC72 catalysts.

Recently, a novel synthesis route has been developed to prepare the bimetallic Pt-Fe/PPycarbon black catalyst [62]. Ultra-fine Pt-Fe dispersion on the PPy-carbon black was obtained by in situ chemical polymerization and co-deposition of Pt and Fe whose ferrous precursor simultaneously functioned as an oxidant for the polymerization of pyrrole. By this novel synthesis route, the metal content in the bimetallic catalysts can be easily controlled. In addition, the process of washing and drying for the preparation of PPy-carbon composite can be omitted, which is a necessary step by a different route [61, 63]. With the in situ interfacial

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chemical polymerization and co-deposition technique, ultra-fine Pt-Fe particles with the size ranging from 2.5 to 3.5nm were dispersed on the PPy-XC72 composite support. The XRD results in Figure 15 indicate that both commercial and Pt-Fe/PPy-XC72 catalyst present a typical fcc crystallographic structure. Compared to the commercial 40% Pt/C catalyst, there is a slight shift of the X-ray diffraction peaks of 25%Pt-10%Fe/PPy-carbon catalyst to higher Bragg angles. It means that some bimetallic interaction or alloying occurs in the catalyst. Two small diffraction peaks at 32.1° and 53.2° belong to Pt3Fe alloy [64]. The inset in Figure 15 illustrates the peak separation of Pt (220) and Pt3Fe (220). Figure 16 shows X-ray photoelectron spectroscopy spectrum of 25%Pt-10%Fe/PPyXC72. The peak at binding energy of 707eV corresponds to metallic iron (Fe 2p3/2) in this catalyst. There are no obvious peaks in the range from 700.4eV to 725eV, which are related to Fe (II) or Fe (III) compounds. The fact that active iron exists in metallic and not in oxidation state suggests that most of the iron may have alloyed or incorporated with platinum, as indicated by XRD result. As for O1s, it should come from the oxidation state of PPy or oxygen-containing groups on carbon surface, such as -COO-.

Figure 16. X-ray photoelectron spectroscopy spectrum of 25%Pt-10%Fe/PPy-XC72 catalyst.

Although the 25%Pt-10%Fe/PPy-XC72 catalyst has a lower Pt loading in the CV measurements, the charge amount of hydrogen adsorption region (-0.25~0.2V vs. Ag/AgCl) is quite similar to that of the commercial 40% Pt/C catalysts (see Figure 17 (I)). The electrochemical active surface area of commercial 40%Pt/C and 25%Pt-10%Fe/PPy-XC72 catalysts is, respectively, about 70 m2 g-1 and 116 m2 g-1, determined from the hydrogen adsorption/desorption region of the CV curves in 0.5M H2SO4 solution. The onset potential for methanol oxidation on 25% Pt-10% Fe/PPy-XC72 was about 0.17 V vs Ag/AgCl (see Figure 17 (II), which is lower than that of the commercial 40% Pt/C (0.21V). In addition, it is evident that the forward anodic peak current density of the 25%Pt-10% Fe/PPy-XC72 is 2 times higher than that of the Pt/C. Therefore, the 25% Pt-10% Fe/PPy-XC72 catalyst possesses a higher specific activity than the 40%Pt/C for methanol oxidation. The high methanol oxidation ability of Pt-Fe/PPy-carbon black implies a potential of the catalyst used as DMFCs anode catalyst.

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In order to reduce the cost of catalyst used in fuel cell, researchers from Los Alamos National Laboratory in USA have discovered a class of non-precious metal composite catalysts using PPy-carbon black (Vulcan XC-72) as catalyst support that could act as PEMFC catalysts [65]. Precious metal catalysts are used to aid fuel oxidation and oxygen reduction between 80-100oC. The reduction reaction at the cathode is considerably slower than the oxidation of H2 at the anode, and requires more catalyst. Efforts to find alternative catalytic materials have therefore focused on non-Pt cathodes, in particular on pyrolized metal porphyrins. These materials demonstrate promising oxygen reduction activity, but lack the stability under fuel cell operating conditions. Bashyam and Zelenay propose a new class of catalysts for O2 reduction that can be synthesized without pyrolysis. Instead, conducting polymer polypyrrole is used as a matrix to entrap transition metal atoms (cobalt) and generate active sites for oxygen reduction.

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(I)

(II)

Figure 17. Cyclic voltammograms of (a) commercial 40% Pt/C (Pt loading of 80 μg cm-2) and (b) 25%Pt-10%Fe/PPy-XC72 (Pt loading of 50 μg cm-2). Electrolyte: (I) 0.5M H2SO4 and (II) 0.5M H2SO4 + 1.0M CH3OH. Scan rate: 50mV s-1, temperature: 25oC.

A schematic diagram of the polypyrrole structure and presumed configuration of the CoPPy is shown in Figure 18. It can be seen that there is a linkage between polypyrrole and Co atoms via nitrogen atoms in the pyrrole units. To ensure electronic conductivity in the catalyst phase, PPy is first synthesized on carbon black (Vulcan XC 72) by in situ chemical polymerization using 10% H2O2 as oxidant and glacial acetic acid as dopant. The performance of the Co/PPy-XC72 nanocompsite catalyst in an H2-O2 PEMFC compares favorably with other non-Pt cathode catalysts, even precious metal-containing alloys, such as Pd-Co-Au and Pd-Ti alloys [66]. The novel catalyst also performs well in a H2-air PEMFC. The performance of the H2-air PEMFC with Co/PPy-XC72 (the Co loading of 6.0×10-7mg cm-2) catalyst in the cathode is very stable, showing no appreciable drop over a 100h period when operated in the range 0.4-0.7V. It opens up the possibility of utilizing a variety of other non-precious composite materials from this class of composites for catalysts of the PEMFC cathode.

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Co-polypyrrole Polypyrrole

C H N Co

Figure 18. Schematic diagram of the Co-polypyrrole composite catalyst.

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Conducting Polymers-Carbon Nanotube Composites Carbon nanotube materials including single-walled carbon nanotubes (SWNT) and multiwalled carbon nanotubes (MWNT) are considered to have several advantages over other conventional catalyst supports for fuel cell applications, which include (i) having more defined crystalline structure with higher electronic conductivity, (ii) containing little impurities, such as metals and sulfides, and thus eliminating potential poisoning effects on catalysts, and (iii) possessing three-dimensional structure thereby favoring the flow of reactants and providing a large reaction zone when fabricated into electrodes. The carbon nanotubes are also chemically stable and resistant to thermal decomposition. Due to these distinctive characteristics, the carbon nanotube material is incredibly suitable for a catalyst support for fuel cell application. However, there are difficulties in dispersing metal nanoparticles at the CNT surfaces. To obtain good nanoparticle mono-dispersion, the surface of CNT must be tailored via proper functionalization. Conducting polymers have proven the suitable host matrices for dispersing metallic particles mentioned above. The composites of conducting polymer with metal nanoparticles permit a facile flow of electronic charges through the polymer matrix during electrochemical processes. In addition, conducting polymers provide a low ohmic drop to the bulk and promote mass transfer for the electroactive species to reach the catalytic sites. Also, metallic particles could be easily dispersed into the matrix of these polymers. Hence, new attempts [67-76] have been made to design and synthesize conducting polymer-carbon nanotubes composites to improve the properties of catalyst materials for fuel cell application. Generally, methods of composite preparation from carbon nanotubes and conducting polymers can be divided into two groups: (1) a mere mixing of conducing polymers and carbon nanotubes, and (2) the composite formation during the polymerization of conducting polymer in the presence of carbon nanotubes. By contrast with the second one, the first type

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of composite shows high resistance due to the difference in the physical characterization between two kinds of different materials, which is not suitable for catalyst support for fuel cell application. The conducing polymer-carbon nanotube composites used as support materials for fuel cell are normally prepared by the electrochemical polymerization of conducting polymer monomers in the presence of carbon nanotubes in electrolyte [67-69, 74, 76] or the chemical polymerization of conducting polymers in the presence of different oxidants and carbon nanotubes [70-73, 75, 76]. Lee et al [68] prepared a novel catalyst for methanol oxidation by dispersing gold nanoparticles onto a PAn grafted multiwall carbon nanotube (MWNT-g-PAn) support through an electrochemical synthesized process. First is to prepare the amine-functionalized MWNT (MWNT-NH2) shown in Figure 19. Then PAn chains were grafted onto MWNT-NH2 by polymerizing a mixture of aniline monomer and MWNT-NH2 dispersed in cetyltrimethyl ammonium bromide via cyclic voltammetric technique. Second is to disperse Au nanoparticles into the film of MWNT-g-PAn by electrochemical reduction of HAuCl4. Au particles of 8-10 nm in size are uniformly distributed into MWNT-g-PAn matrix. Au/MWNTg-PAn catalysts exhibit excellent catalytic activity towards the oxidation of methanol and COads. Methanol oxidation occurs at a much lower oxidation potential (780 mV) at the Au/MWNT-g-PAn electrode with high current density compared with the Au/MWNT electrode (890 mV) and pristine Au electrode (930 mV), due to the dispersion of Au nanoparticles into the three-dimensional network of MWNT-g-PAn. Moreover, the Au/MWNT-g-PAn catalyst has better oxidation kinetics for the oxidation of COads to CO2 than the Au/MWNT and pristine Au electrodes. (A) Functionalization of MWNT

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HNO3

SO2Cl

NH2-R-Y

65oC, 24 h

60oC, 24 h

MWNT

R: PEG, Y: NH2

(B) Fabrication of PAn-g-MWNT-Au electrode Au

Aniline monomer + MWNT-NH2

MWNT-g-PAn

Electrochemical

HAuCl4

Polymerization

Electrochemical reduction

MWNT-g-PAn

MWNT-g-PAn-Au Au nanoparticle

Figure 19. Schematic diagram of fabrication of Au/MWNT-g-PAn catalyst.

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Very recently, Li et al developed a novel method to prepare homogeneous PAn-carbon nanotubes composites by electrochemical codeposition as catalyst support [74, 76]. The procedure for the preparation of Pt/PAn-MWNT catalysts is shown in Figure 20. The first step is to functionalize the MWNT via the diazotization reaction. The 4-carboxylicbenzene groups are modified on the surface of MWNT via C-C covalent bonds which are strong and propitious to disperse them well in the aniline. Then the functional MWNTs are added to aniline monomer solution and heated at reflux for 3 h. Whereafter, the aniline monomer in 1.0M H2SO4 solution is electrochemically polymerized using cyclic voltammetric method. Finally, Pt is dispersed on the PAn-MWNT composite by electrodeposition technique. Since the MWNT in the composite is covalently functionalized with PAn, it can ensure the compatibility of carbon nanotubes in the PAn matrix, which can avoid potential microscopic phase separation in the nanocomposite. Furthermore, the incorporation of MWNT also leads to more active sites for faradic reaction and higher accessible surface area than pure PAn. Meanwhile, the special frame and properties of the PAn-MWNT composites help to achieve a fine dispersion of Pt particles in the composites. Therefore, Pt/PAn-MWNT catalysts show higher catalytic activity and better long-term stability than Pt/PAn catalysts towards formic acid oxidation. In addition, some researchers also found that PAn modified SWNT or MWNT composites, prepared by electrochemical polymerization of aniline containing well-dispersed carbon nanotubes, were also promising catalyst supports for methanol oxidation [67, 69].

Figure 20. Schematic procedures for the preparation of Pt-modified PAn-MWNT composite films.

Conducting PAn and PPy polymers modified with carbon nanotubes have been prepared by simple but efficient chemical polymerization of aniline [70] or pyrrole [71-73, 75] monomers on the surface of carbon nanotubes. The catalytic activities and stabilities of Pt, PtRu and Pt-Pd deposited on the PPy-MWNT supports towards the oxidation of methanol, formic acid and formaldehyde were investigated by Alagar et al [71, 72]. The composite catalysts were prepared by in situ polymerization of pyrrole on carbon nanotubes in HCl aqueous solution containing (NH4)S2O8 as oxidant. A recent paper describes cobalt-PPy-

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MWNT composite as an catalyst for the oxygen reduction in PEMFC [73]. Uniform dispersion of PPy over MWNT has been achieved by in situ chemical polymerization technique via oxidizing monomer using FeCl3 as oxidant, and nanosized particles of Co can be dispersed on PPy-MWNT composite as shown in Figure 21. The suitability of this catalyst for the oxygen reduction in direct methanol (or ethanol) fuel cell has been examined. The maximum power density of 140mW cm-2 is obtained with the catalyst loading of 0.25mg cm-2 of Pt-Ru and 0.5mg cm-2 of Co on the anode and cathode, respectively (Figure 22). In the same operated conditions, PEMFC generates 140mW cm-2 with Pt/C anode and Pt/C cathode at the same catalyst loading. This means that Co/PPy-MWNT catalyst has higher oxygen reduction activity than the Pt/C catalyst. In addition, Co/PPy-MWNT catalyst in PEMFC is quite stable for 50h not only at high current density but also at high voltage (0.7V). The higher oxygen reduction activity of non-precious Co/PPy-MWNT at the cathode electrode than Pt/C has also been demonstrated in direct alcohol fuel cell. These results show that the cobalt-PPy-MWNT composite is a promising candidate catalyst for the oxygen reduction in fuel cells.

Figure 21. TEM (A) and HR-TEM (B) images of Co/PPy-MWNT catalyst [73].

Figure 22. Polarization curves of PEMFC at various temperature with Co/PPy-MWNT (loading: 0.5mg cm-2) and Pt-Ru/MWNT (loading: 0.25mg cm-2) as cathode and anode catalysts, respectively [73].

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We have noticed that the catalytic activity of metal catalysts on the polypyrrole modified carbon nanotubes towards the oxidation of hydrogen and methanol as well as reduction of oxygen can be significantly enhanced after over-oxidation treatment. The over-oxidation treatment is to extend the potential window to an upper limit of 1.4V versus Ag/AgCl electrode. The process is carried out in a cyclic potential range from -0.3 to 1.4V at 50mV s-1 sweep rate in 1.0M H2SO4 solution. Under these conditions, irreversible over-oxidation of PPy occurs. In order to check the effect of the over-oxidation treatment to the catalyst structure, the untreated and treated catalysts were ultrasonically dispersed in ethanol solution, and coated on ultrathin carbon membrane before TEM measurements. Typical TEM images of catalysts before and after over-oxidation treatment are exhibited in Figure 23. The over-oxidation activation process does not affect the particle size obviously. The lattice distance of Pt (111) space calculated from the selecting area electron diffraction patterns for the untreated and treated catalysts is 2.21Å and 2.17Å, respectively. It can be observed from Figure 24 that PPy geometric configuration changes, and more Pt particles bunched in PPy matrix are exposed after the over-oxidation process. The reason is that PPy is oxidized from a reduction state to an oxidant state during the over-oxidation process, which will break the long chains of PPy into the short chains and make more Pt particles expose.

Figure 23. TEM images of Pt3Co/PPy-MWNT catalysts before (a) and after (b) over-oxidation treatments.

Figure 24. SEM morphologies of Pt3Co/PPy-MWNT catalysts (a) before and (b) after over-oxidation treatments.

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Before the over-oxidation treatment, the characteristic adsorption and desorption peaks of hydrogen oxidation and oxygen reduction which usually appear on pure Pt electrode are not obvious, which means that Pt/PPy-MWNT catalysts have low catalytic acitivity. The reason is that most of Pt particles have been bundled and embedded in PPy network and are not easy to contact with H2 and O2. During over-oxidation in the range of -0.3 to 1.4V about 50 cycles as shown in Figure 25a, the characteristic Pt peaks become more and more prominent in the initial cycles and then stable. A comparison of the CV curves of Pt3Co/PPy-MWNT in the range of -0.3 to 1.0V (Figure 25b) indicates that the catalytic activity of catalysts is improved obviously after the over-oxidation activation process. A clear enhancement in the catalytic activity to methanol oxidation can be also observed from Figure 26. Moreover, we find that this activation method is also valid for Pt/PPy or Pt/PPy-MWNT catalysts. All these positive results can be explained by the fact that the over-oxidation treatment leads to the partial degradation of PPy chains, an increase in porosity of the catalyst support matrix, and gas accessibility to the active centers for electrochemical reaction.

Figure 25. Cyclic voltammograms of Pt3Co/PPy-MWNT catalysts in 0.5M H2SO4 at 50mV s-1 and 25oC. (a) the first 50 cycles during the over-oxidation treatment, (b) a comparison of the catalyst activity before and after over-oxidation treatments. Pt loading: 500μg cm-2.

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Figure 26. Anodic scanning curves of Pt3Co/PPy-MWNT catalysts in 0.5M H2SO4 + 1.0M CH3OH at 50mV s-1 and 25oC. (a) before and (b) after over-oxidation treatment. Pt loading: 500μg cm-2.

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Other Conducting Polymer Composites Conducting polymers incorporated with minerals have been developed for catalyst supports of fuel cells. Rajesh et al [77] prepared transition metal oxide (vanadium pentoxide) and PAn composite as catalyst support for Pt nanoparticles. The aniline monomer is in situ chemically polymerized by V2O5, along with the generation of V4+. V2O5 not only acts as the host lattice for the intercalation of polyaniline, but also as an oxidant in the polymerization process. V2O5 in Pt/PAn-V2O5 catalysts will play a role in effectively removing the nascent hydrogen from the Pt surface during dehydrogenation of methanol. This helps to perform the dehydrogenation over Pt at relatively low over-potentials. Also the oxophilic nature of vanadium pentoxide helps in removing the strongly adsorbed reaction intermediates from the reaction sites. Pt nanoparticles dispersed on PAn-V2O5 support have exhibited excellent catalytic activity and stability for methanol oxidation compared to Pt dispersed on the commercial Vulcan XC-72R carbon black. This is, at least in part, related to the better CO tolerance of the PAn-V2O5 composite materials. PAn not only acts as an excellent electronic support for Pt particles, but also ensures the stability of the layered transition metal oxide under the electrochemical operating conditions. Since carbon material is impermeable to gases and liquid, it would make the liquid alcohol hardly get into the catalyst layer. It not only reduces the three-phase interface, but also decreases the electrode kinetics, resulting in the lower catalyst utilization and higher cost of fuel cell systems. In order to solve this problem, a novel three-dimensional electrode has been developed using polypyrrole coated polystyrene sphere (PS) covered by Pt nanoparticles [78, 79]. The three-dimensional electrode was fabricated by the following three steps: (1) the preparation of PS with 200nm-2µm by emulsifier-free-dispersion polymerization method, (2) the in situ chemical polymerization of PPy on the surface of PS to make the support electrically conductive, (3) the deposition of catalyst on the surface of PPy and the fabrication of three-dimensional electrode by coating the catalyst ink on the conducting substrate. Such a structure permits liquid alcohol to diffuse into the catalyst layer easily and form larger threephase interface for electrochemical reactions. In fact, an improved performance for methanol

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oxidation on this novel electrode has been obtained compared to a conventional electrode made by commercially available E-TEK catalyst with the same Pt loading. This new type of electrode is promising for liquid fuel cells.

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[47] Bensebaa, F., Farah, A.A., Wang, D., Bock, C., Du, X., Kung, J., & Page, Y.L. (2005). Microwave synthesis of polymer-embedded Pt-Ru catalyst for direct methanol fuel cell, J. Phys. Chem. B, 109, 15339-15344. [48] Kim, S., & Park, S.-J. (2008). Electroactivity of Pt–Ru/polyaniline composite catalystelectrodes prepared by electrochemical deposition methods. Solid State Ionics, 178, 1915-1921. [49] Jüttner, K., Mangold, K.-M., Lange, M., & Bouzek, K. (2004). Preparation and properties of composite polypyrrole/Pt catalyst systems. Russ. J. Electrochem., 40, 317325. [50] Qi, Z., & Pickup, P.G. (1997). Size control of polypyrrole particles. Chem. Mater., 9, 2934-2939. [51] Qi, Z., & Pickup, P.G. (1998). Novel supported catalysts: platinum and platinum oxide nanoparticles dispersed on polypyrrole/polystyrenesulfonate particles. Chem. Commun., 15-16. [52] O’Mullane, A.P., Dale, S.E., Macpherson, J.V., & Unwin, P.R. (2004). Fabrication and electrocatalytic properties of polyaniline/Pt nanoparticle composites. Chem. Commun., 1606-1607. [53] MacDiarmid, A.G., Manohar, S.K., Masters, J.G., Sun, Y., Weiss H., & Epstein, A.J. (1991). Polyaniline: synthesis and properties of pernigraniline base. Synth. Met., 41, 621-626. [54] Lei Li unpublished results [55] Michel, M., Ettingshausen, F., Scheiba, F., Wolz A., & Roth, C. (2008). Using layerby-layer assembly of polyaniline fibers in the fast preparation of high performance fuel cell nanostructured membrane electrodes. Phys. Chem. Chem. Phys., 10, 3796-3801. [56] Li, L., Zhang, Y., Drillet, J.-F., Dittmeyer, R., & Jüttner, K. (2007). Preparation and characterization of Pt direct deposition on polypyrrole modified Nafion composite membranes for direct methanol fuel cell applications. Chem. Eng. J., 133, 113-119. [57] Choi, J.-H., Park, K.-Y., Lee, H.-K., Kim, Y.-M., Lee, J.-S., & Sung, Y.-E. (2003). Nano-composite of PtRu alloy electrocatalyst and electronically conducting polymer for use as the anode in a direct methanol fuel cell. Electrochim. Acta, 48, 2781-2789. [58] Roy, S.C., Christensen, P.A., Hamnett, A., Thomas, K.M., & Trapp, V. (1996). Direct methanol fuel cell cathodes with sulfur and nitrogen-based carbon functionality. J. Electrochem. Soc., 143, 3073-3079. [59] Wu, G., Li, L., Li, J., & Xu, B. (2005). Polyaniline-carbon composite films as supports of Pt and PtRu particles for methanol electrooxidation. Carbon, 43, 2579-2587. [60] Ozenler, S.S., & Kadırgan, F. (2006). The effect of the matrix on the electro-catalytic properties of methanol tolerant oxygen reduction catalysts based on rutheniumchalcogenides. J. Power Sources, 154, 364-369. [61] Zhao, H., Li, L., Yang, J., & Zhang, Y. (2008). Nanostructured polypyrrole/carbon composite as Pt catalyst support for fuel cell applications. J. Power Sources, 184, 375380. [62] Zhao, H., Li, L., Yang, J., Zhang, Y., & Li, H. (2008). Synthesis and characterization of bimetallic Pt–Fe/polypyrrole–carbon catalyst as DMFC anode catalyst. Electrochem. Commun., 10, 876-879.

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[63] Xu, Y., Peng, X., Zeng, H., Dai, L., & Wu, H. (2008). Study of an anti-poisoning catalyst for methanol electro-oxidation based on PAn-C composite carriers. C. R. Chimie, 11, 147-151. [64] Cabri, L.J., & Feather, C.E. (1975). Platinum-iron alloys: a nomenclature based on a study of natural and synthetic alloys. Can. Mineral., 13, 117-126. [65] Bashyam, R., & Zelenay, P. (2006). A class of non-precious metal composite catalysts for fuel cells. Nature, 443, 63-66. [66] Fernandez, J.L., Raghuveer, V., Manthiram, A., & Bard, A.J. (2005). Pd-Ti and Pd-CoAu electrocatalysts as a replacement for platinum for oxygen reduction in proton exchange membrane fuel cells. J. Am. Chem. Soc., 127, 13100-13101. [67] Shi, J., Wang, Z., & Li, H. (2007). Electrochemical fabrication of polyaniline/multiwalled carbon nanotube composite films for electrooxidation of methanol. J. Mater. Sci., 42, 539-544. [68] Santhosh, P., Gopalan, A., & Lee, K.-P. (2006). Gold nanoparticles dispersed polyaniline grafted multiwall carbon nanotubes as newer electrocatalysts: preparation and performances for methanol oxidation. J. Catal., 238, 177-185. [69] Wu, G., Li, L., Li, J., & Xu, B. (2006). Methanol electrooxidation on Pt particles dispersed into PANI/SWNT composite films. J. Power Sources, 155, 118-127. [70] Konyushenko, E.N., Stejskal, J., Hradil, M.T.J., Kovářová, J., Prokeš, J., Cieslar, M., Hwang, J.-Y., Chen, K.-H., & Sapurina, I. (2006). Multi-wall carbon nanotubes coated with polyaniline. Polymer, 47, 5715-5723. [71] Selvaraj, V., & Alagar, M. (2007). Pt and Pt–Ru nanoparticles decorated polypyrrole/multiwalled carbon nanotubes and their catalytic activity towards methanol oxidation. Electrochem. Commun., 9, 1145-1153. [72] Selvaraj, V., Alagar, M., & Kumar, K. S. (2007). Synthesis and characterization of metal nanoparticles-decorated PPY–CNT composite and their electrocatalytic oxidation of formic acid and formaldehyde for fuel cell applications. Appl. Catal. B-Environ., 75, 129-138. [73] Reddy, A.L.M., Rajalakshmi, N., & Ramaprabhu, S. (2008). Cobalt-polypyrrolemultiwalled carbon nanotube catalysts for hydrogen and alcohol fuel cells. Carbon, 46, 2-11. [74] Zhu, Z., Wang, Z., & Li, H. (2008). Functional multi-walled carbon nanotube/polyaniline composite films as supports of platinum for formic acid electrooxidation. Appl. Surf. Sci., 254, 2934-2940. [75] Kim, J.-Y., Kim, K.H., & Kim, K.B. (2008). Fabrication and electrochemical properties of carbon nanotube/polypyrrole composite film electrodes with controlled pore size. J. Power Sources, 176, 396-402. [76] Wang, Z., Zhu, Z., Shi, J., & Li, H. (2007). Electrocatalytic oxidation of formaldehyde on platinum well-dispersed into single-wall carbon nanotube/polyaniline composite film. Appl. Surf. Sci., 253, 8811-8817. [77] Rajesh, B., Thampi, K.R., Bonard, J.-M., Mathieu, H.J., Xanthopoulos, N., & Viswanathan, B. (2005). Electronically conducting hybrid material as high performance catalyst support for electrocatalytic application. J. Power Sources, 141, 35-38. [78] Xie, F., Tian, Z., Meng, H., & Shen, P. (2005). Increasing the three-phase boundary by a novel three-dimensional electrode. J. Power Sources, 141, 211-215.

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[79] Xie, F., Meng, H., & Shen, P. (2008). Diffusion study in a novel three-dimensional electrode for direct methanol fuel cells. Electrochim. Acta, 53, 5039-5044.

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In: Polymer Electrolyte Membrane Fuel Cells… Editors: R. Esposito, A. Conti

ISBN: 978-1-60692-773-1 ©2009 Nova Science Publishers, Inc.

Chapter 12

ELECTROCATALYTIC MODIFIED ELECTRODES WITH TRANSITION METAL AZAMACROCYCLES AND OTHER COMPLEXES FOR THE DETECTION OF SULFUR AND NITROGEN OXOANIONS Maria J. Aguirre1, Yo-ying Chen1, Galo Ramirez2, Mauricio Isaacs3, Fabiola Isaacs1 and William Cheuquepan1

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1

Departmento of Química de los Materiales, Facultad de Química y Biología Universidad de Santiago de Chile, USACH. Av. B. O’Higgins 3363, Estación Central, Santiago, Chile. 2 Universidad Andres Bello, Departamento de Ciencias Químicas, Facultad de Ecología y Recursos Naturales, República 275, Santiago, Chile. 3 Departmento de Química, Facultad de Ciencias, Universidad de Chile, UCHILE, Las Palmeras 3425, Ñuñoa, Santiago, Chile.

ABSTRACT Modified electrodes with transition metal phenantrolines, phthalocyanines, porphyrins and other complexes where the metal center is a transition metal have been widely studied because they show highly activity toward the oxidation and reduction of several analytes. One of the most interesting applications of these modified electrodes is the determination of oxoanions in waste water, food and beverages. In fact, they have been used to detect nitrite, nitrate, sulfite, and other pollutants or traces with low limit detection and in some cases; specificity. On the other hand, it is interesting that these modified electrodes can be designed for desirable purposes by changing the metal center, modifying the ligand, or changing the electrolyte. There are some cases in which the position of a substituent on the ligand can drastically change its activity. The modified electrodes are been used as supramolecular assemblies, as electroactive films or as composites of different layers adsorbed on the surface. This brief review shows the last decade’s studies on the detection of sulfur and nitrogen oxoanions using electrodes modified with metal transition complexes.

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I. NITROGEN OXOAANIONS Nitrate and nitrite are omnipresent within environmental, industrial and physiological fields, and while the understanding of their role within such matrices has increased, a substantial degree of doubt and conjecture remains. These ions have been commercially exploited throughout the ages to promote the development of mankind but there is no doubt that our affection for them has declined in recent years. The constant use of, and indeed reliance upon, these versatile agents combined with revelations of their potential toxicity have raised numerous concerns [1-3]. These problems have been widely recognised, and, as a result, legislative frameworks aimed at controlling their level within the wider environment and within food products have been imposed in most industrialised countries [1, 4, 5]. The restrictions placed upon the commercial utilisation of these ions have eased some of the apprehensions raised by the medical community but the establishment of adequate controls can only be achieved by fully exposing their influence on the various pathways that govern environmental and physiological well being. The need and desire to monitor these ions are undeniable, yet their ubiquity can pose a significant challenge to the analytical community. Nitrate and nitrite have become intertwined with domestic life, and it is effectively impossible to engage in everyday activities without encountering these ions or the products of their use. The chemical versatility of these anions has ensured their use within a multitude of industrial processes ranging from the manufacture of fireworks to the production of the latest dyes. Their antimicrobial action has been recognised for centuries and is still used for the preservation of meat products. Despite the huge number of products that are dependent upon these ions, it is their association within environmental issues that has captured the interest of the public and a significant proportion of the scientific community. Inputs of nitrate and nitrite to the environment can occur through industrial and domestic combustion processes with gaseous NOx species converted to NO3− through photochemical conversion within the atmosphere [1, 6–10]. The vast majority, however, arise from agricultural sources [1,11]. The mishandling of inorganic fertilisers combined with the more general mismanagement of the natural resources has been suggested as resulting in the perturbation of both local and global nitrogen cycles [1, 11, 12]. The consequences of the environmental manipulations are as yet uncertain, and therefore monitoring the ecological fate of nitrate and nitrite has gained increasing importance. The high solubility and mobility of these ions within the soil and the continued reliance on inorganic fertiliser has led to reports that ‘‘run off’’ is a continual hazard wherever agricultural processes are in close proximity to surface water [1, 13]. Eutrophication of lakes and more recently coastal outfalls are thought to result in the generation of algal blooms that wreak disaster with local ecological systems [1, 14, 15]. The contamination of edible shellfish and the occurrence of ‘‘red tides’’ of potentially toxic algae near tourist resorts can also impart a degree of economic misery to the afflicted communities [1, 16]. The potential contamination of groundwater through the percolation of nitrates through natural aquifers presents the most immediate risk to health [1-3] and as such, the maximum permissible level for these ions in drinking water supplies is often levied and currently stands at 50 mg/L in the UK [1, 4, 5, 17]. The two main threats to health that arise from the ingestion of these ions are reported as ‘‘blue baby’’ syndrome and gastric cancer [1-3, 18, 19]. In both cases, the principal protagonist is nitrite obtained directly from contaminated water supplies

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or derived from the reduction of nitrate by the multifarious bacterial colonies that reside within the mouth. Passage of nitrite into the bloodstream results in the irreversible conversion of hemogloblin to methemoglobin with oxygen uptake and transportation compromised [1, 3]. This is particularly hazardous for infants, given their limited physical stature and the susceptibility of their neural development to impeded oxygen transport. A more contentious health concern is the possible formation of carcinogenic nitrosoamines within the acidic conditions of the stomach and their subsequent implication in the pathology of gastric cancer [1, 2, 20]. Upon reaching the stomach, nitrite is converted to nitrous acid, which can act as a powerful nitrosating agent. While nitrosoamines can be carcinogenic, conclusive evidence linking nitrite ingestion with the stomach cancer remains obscure [1, 3]. The presence of nitrate and nitrite within physiological systems is often viewed with considerable and, as noted above, justifiable concern. Ingestion is not, however, the sole source by which these ions can arise within physiological systems. The endogenous production of nitrate and nitrite within tissue can occur as a result of the activity of their more transient cousin, nitric oxide [1, 21, 22]. The reaction of nitric oxide with oxygen leads to the production of nitrite [1, 2], and it is through the greater stability of this ion that the action of NO takes place:

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4 NO + O2 + 2H2O → 4NO2− +4 H+

Nitric oxide has been shown to play an important role in many metabolic functions, including the regulation of vascular tone, inhibition of platelet aggregation, neurotransmitter, cytotoxic agent, thrombosis, and inflammation, and may also play a role in the immune system [1, 23–25]. The measurement of nitrite can therefore provide a reliable measurement of NO action within the body, and, as such, can be used as a biomarker that enables physicians to gauge the health of an individual [1, 26]. This has particular importance when assessing inflammatory processes as the level of nitrite can be correlated to the degree of injury. Among the more common aliments for which the monitoring of nitrate and nitrite can be beneficial are sepsis [1, 27], infectious gastroenteritis [1, 28], meningitis [1, 29], Parkinson’s disease [1, 20], minimal change nephritic syndrome in children [1, 30], preeclampsia [1, 31] and rheumatoid arthritis [1, 32]. Analytical methods for determination of nitrate are based in several techniques such as chemiluminiscence [1, 8] IR [1, 33] Raman [34] and electrochemical detection [1, 35-37], but the most used method it is an indirect one based in the Greiss’s reaction that involves the previous conversion to NO2- by several reductant agents (copperized Cd columns for instance). The assay typically relies on the diazotisation of a suitable aromatic amine by acidified nitrite, with the subsequent coupling reaction providing a highly colored azo chromophore from which the concentration of nitrite can be determined [1]. Coordination chemistry focused in macrocyclic ligands containing transition metals appears as an opportunity to develop electroanalytical sensors, since macrocyclic ligands are capable to form stable complexes with a great variety of metals. Depending on the substituents in the periphery of the ring, the metallic central ion inserted in the cavity, (Mn (II), Co (II), Fe (II) or Ni (II)) and after coordination of the analyte in the 5th or 6th axial position, the reaction under study it can show a high catalytic efficiency, due to the rich redox chemistry provided by these systems. All these properties can be used to design modified electrodes that eventually can use in the direct determination of several target molecules,

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NO3- in this case. However the literature about direct electrochemical reduction of NO3mediated by macrocyclic coordination compounds is scarce, compared with similar N oxo compounds such as NO and, NO2-, probably due to the high stability of the anion where the central atom is in last oxidation state (5+). This part of the chapter will be focused in the description of the electrochemical reduction of NO3- by cyclams and related macrocycles, on several electrodic surfaces and its application as potentiometric sensors in the NO3- determination.

I. 1. Nitrate Determination

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Co (III) cyclam (cyclam = 1,4,8,11-tetraazacyclotetradecane) (see Structure 1) has been demonstrated to act as a catalyst for nitrate reduction by cyclic voltammetry and controlledpotential electrolysis (CPE) [38].

Structure 1. M= Co or Ni.

CPE of a KNO3 solution at an Hg electrode gave hydroxylamine selectively with a high current yields (ca. 90%); the current densities were > 5mA / cm2 in 0.1 M KNO3 solution at 1.5 V (vs SCE) in the presence of 20 μM or less of the catalyst. The turnover frequency was > 500 h-1 (using an 8 cm2 Hg pool electrode), and no deactivation of the catalyst was observed after CPE for several hours. At Ag, Cu and Pb electrodes, hydroxylamine was reduced further to give ammonia preferentially. Ni(II)-cyclam also acted as a catalyst, but the effective potential region was ca. 0.2 V more negative than that of Co(III)-cyclam. UV-vis experiments suggested that Co(III)- and Ni(II)-cyclam did not bind to NO3- upon the addition up to 1M because not changes in the respective spectra were observed. Instead the formation of Co(I) and Ni(I)-cyclam, which adsorbed onto the electrode surface, would be effective in the electrocatalytic pathway. Also the role of cyclam as a ligand would be providing an adequate structure and electronic density in the metal center ion for the interaction with the N oxoanions [38].

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On the other hand Co(III)-cyclam catalyst [39] was incorporated into a Nafion film Au, Au/Hg electrodes, and that the reduction of nitrate proceeds efficiently at these electrodes in concentrated NaOH solution, in this case the incorporation of this complex in the perflouroalkanesulfonated membrane enhanced the potential at which the reduction of nitrate takes place by about 100 mV. The addition of low concentrations of CrO4- does not change the voltammetric response of these modified electrodes suggesting some degree of selectivity. Mechanistic insight of these modified electrodes were obtained when Co(III)-cyclam incorporated into a Nafion film Au disk electrodes were employed to catalyze the reduction of nitrate in concentrated NaOH solution [40]. The kinetic characteristics of this catalytic system were determined by the techniques of rotating-disk electrode voltammetry and cyclic voltammetry with the aid of the developed kinetic models. The rate constant of the electrocatalytic reaction was ca.180 Ms with no dependence on the interfacial concentration of the complex on the electrode. Also the catalytic reduction on this modified electrode was controlled by a mixture of nitrate diffusion in the film and the catalytic reaction itself [40]. The authors proposed that in the presence Co(I)-(cyclam) NO3, as a key mechanistic path, the intermediary is generated in the diffusion layer at ca. -1.3 V vs SCE. The lack of dependence on pH of either the catalytic peak potential or the peak current indicates that the metal hydride intermediate species are not involved in the initial steps of the reduction process. Both Co(I) and Co(III) intermediates in the nitrate reduction were detected in collection experiments with a gold ring-disk electrode [41]. In this case a mixture of products including hydroxylamine and ammonia was obtained. On the other hand the use of Co-DIM (DIM= 2, 3 dimethyl-1,4,8,11 tetraazacyclotetradeca-1,3-diene) a Co-cyclam analog (see Structure 2), was reported to catalyze the electrochemical reduction of nitrate to ammonia [42].

N

N

CH3

N

CH3

M N

Structure 2. M=Co.

The Co(II)/Co(I) reduction of this compound in aqueous bromide solutions was nearly reversible on pyrolytic graphite electrodes. The complex was catalytically active for the reduction of NO3- as well as its reduction intermediates NO2- and hydroxylamine. Ammonia was the sole product found after electrolysis of sodium nitrate in the presence of Co-DIM on a carbon felt electrode. Another related study showed the N- functionalization of cyclam and the linked to the nitrogen of pyrrole through an alkyl spacer (see structure 3) [43].

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H N

N

R

M N

N

H

H

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Structure 3. M = Co, Ni, R = -(CH2)5-N-Pyrrol.

After complexation of the ligand by Ni(II) and Co(II), the complexes were electropolymerized on the surface of several electrode surfaces by cyclic voltammetry or at constant potential. The Ni(III)/ Ni(II) redox system is very distinctly visible on the voltammogram of poly[Ni(II)-cyclam-N-(CH2)5-pyrrole] film, even after over-oxidation of the polypirrole backbone; from this observation the authors [43] claimed a hopping electron transfer mechanism between nickel sites. Several differences were observed between the electropolymerization of Co(III) and Co(II) complexes. Whereas the film obtained from Co(III)-cyclam-N-(CH2)5pyrrole were fragile; those obtained from the electropolymerization of the analogous Co(II) complex was of better quality. The Co modified glassy carbon electrodes were active for the electrocatalytic reduction of dioxygen in acid media. A gold electrode modified by poly[Co(III)-cyclam-N-(CH2)5-pyrrole] catalyses the nitrate ion reduction at -1.0 V vs SCE in a strongly basic medium after amalgamation [43]. On the other hand the Ni(II) 5,7,12,14-tetramethyl-1,4,8,11- tetraazacyclotetradeca4,6,11,13-tetraene complex (see Structure 4) [44] was used as neutral carrier in plasticized PVC membrane to test the potentiometric response and properties toward the nitrate determination, where relatively high selectivity was obtained.

H3C

CH3 N

N M

N H3C

N CH3

Structure 4. M= Ni.

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The influence of several variables was investigated in order to optimize the potentiometric response and selectivity of the electrode. The resulting membrane electrode incorporating 31% PVC, 61% dioctyl phpthlate (DOP) as plasticizer, 3% methyltrioctyl ammonium chloride (MTOAC) as cationic additive and 5% carrier (all w/W) demonstrated a Nernstian response slope of 59,6 mV per decade over the concentration range of 5x10-6 – 1x10-1 M NO3-. The electrode exhibits a fast response time (≤ 10 s), a detection limit of 2.5x10-6 M, and can be used over a wide pH range of 4-12. This electrode was successfully applied to the determination of nitrate ion natural water samples [44]. Other example was presented by the same authors. In this case another nitrate-selective coated-wire electrode based on the complex (6,8,15,17-tetramethyl-7H,16H-5,9,14,18-tetraazidobenzo[b,i]-cyclotetra-decanato-(2)-K4-N,N’,N’’,N’’’) Ni (II) (see Structure 5) [45] as the membrane carrier was developed.

H3C

CH3 N

N M

N H3C

N CH3

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Structure 5. M = Ni.

The electrode exhibits a good Nernstian slope of -57.8 mV/decade and a linear range of 1.0 x 10-5- 1.0 M for NO3-. The limit of detection is 5x10-6 M. It has a fast response time of 15 s it was used for more than four months. The selective coefficients were determined by the fixed interference method and could be in the 2.2-9.4 pH range. The electrode was employed as an indicator electrode for direct determination of NO3- in fresh water. To the contrary it seems to be that changing the coordination sphere of this same metal ions the sensing properties also can change to nitrite, for instance metallo-salens of Co(II) (Co-Sal), Cr(III), (Cr-sal) and Al(III) (Al-sal) (see structure 6) [46] were used as the active ionophores within plasticized PVC membranes. It is shown that central metal ion plays a critical role in directing the ionophore selectivity. Polymer-membrane electrodes based on Co-Sal, Cr-Sal, and Al-Sal were demonstrated to exhibit enhanced responses and selectivity toward nitrite /thiocyanate, thiocyanate, and fluoride anions respectively. The anion selectivity of the three coordination compounds ionophores systems is shown to deviate significantly from the classical Hoffmeister pattern that is based only on ion lipophilicity. For example, optimized membrane pot

electrodes for nitrite ion based on Co-Sal exhibit log K Nitrite , Anion values of -5.22, -4.66, 4.48, -2.5 towards bromide, perchlorate, nitrate, and iodide anions, respectively.

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H N

H3C

H N

H3C

M O

H3C

CH3

H3C

CH3

O

CH3

CH3

CH3

H3C

CH3

H3C

CH3

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Structure 6. M = Co (II), Cr (III), Al (III).

Optimized membrane electrodes based on Co-Sal and Cr-Sal show near-Nerstian response toward nitrite [46] (-57mV/decade) and thiocyante (-56.9 mV/ decade), respectively, with fast response and recovery times. In contrast, Al-Sal based membrane electrodes respond to fluoride ion in a super-Nernstian (-70 mV/decade) and nearly irreversible mode. The operative response mechanism of Co-Sal, Cr-Sal and Al-Sal membrane electrodes was examined using the effect of added ionic sites on the potetiometric response characteristics. It was demonstrated that addition of lipophilic anionic sites to membrane electrodes based on the utilized metallo-salens enhanced the selectivity towards the primary ion, while addition of cathodic sites resulted in Hofmeister selectivity patterns suggesting that the operative response mechanism is of the charged carrier type. Electron spin resonance data indicate that Co(II) metal-ion center of Co-Sal complex ionophore undergoes oxidation to Co(III). This process leads to formation of a charged anion-carrier that is consistent with the response behaviour obtained for Co-Sal based membrane electrodes [46]. The same Al-Sal complex also demonstrated to be useful as an optical sensor toward fluoride anion [47]. The sensor was prepared by embedding the aluminum(III)-salen ionophore and a suitable lipophilic pH-sensitive indicator (ETH-7075) in a plasticized poly(vinyl chloride) (PVC) film. Another interesting system of polymeric membranes containing Zn(II) complexes as anion carriers was prepared for determination of nitrate anion present in water samples [48]. Two Zn(II) complexes coordinated by neutral tetradentate ligands, N,N_-ethylene-bis(Nmethylamide) and N,N_-ethylene-bis(N-methyl-(S)-alanine methyl-(S)-alanine dimethylamide), (see structure 7) worked well as anion-selective carriers [48].

O H3C

CH3

CH3 N

N

N H

H

N CH3

CH3

CH3 O

L1

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O H3C

CH3

CH3 N

CH3 N

N CH3

349

N CH3

CH3

CH3 O

L2 Structure 7.

[Zn(L)(CH3OH)n(H2O)2-n](ClO4)2, L=L1,L2

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The combination of these new Zn(II) complexes with dioctylsebacate as a plasticizer particularly offered high sensing selectivity for nitrate anion. They exhibited near-Nernstian slopes in the wide linear concentration range of 5.0×10−5 to 1.0×10−1 M, and operated well in the wide pH range from 4 to 11 with the response time of less than 25 s. The potentiometric selectivity coefficients were evaluated using the fixed interference method, indicating that the two Zn(II) complexes exhibited better selectivity for nitrate anion with respect to a wide variety of inorganic anions. Although chloride anion worked as an interfering species at a concentration higher than 1.0×10−3 M, the Zn(II) complex-based sensors were applicable in determination of the nitrate anion after adding silver sulfate to remove the chloride anion [48]. The amperometric detection of NO3- on a glassy carbon Co(II) tetrarhutenated porphyrin/ Zn(II) tetrasulphonated porphyrin (see structure 8) [49] layer by layer modified electrode was achieve.

Structure 8. R= SO3-..

From the reaction of nitrate after reduction of nitrite in a reductor column containing copperised cadmium. This method was used for the determination of nitrate and nitrite in mineral water, saliva and cured meats. Polymeric-membrane electrodes for NO3− anion based on 5,10,15,20-tetrakis [2-(4methylphenylurea)phenyl]porphyrin and similar urea-functionalized porphyrins as neutral ionophores were prepared (see Structure 9) [50].

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Structure 9. Structure taken from Reference 50 (Lee, H.; Song, K.; Seo, H.; Jeon, S. Talanta 2004 62 293). Reprinted from Talanta, Volume 62(2), H. K. Lee, K. Song, H.R. Seo, S. Jeon, Nitrate-selective electrodes based on meso-tetrakis [((2-arylphenylurea)phenyl]porphyrins as neutral lipophylic ionopheres, Figure 1, Copyright (2008), with permission from Elsevier.

The effects of polymeric-membrane composition and test solution pH were tested. The electrodes revealed good selectivity coefficients for NO3− over a wide variety of other anions, and the values of eight anions were measured. Of the various electrodes prepared, the electrode based on urea-functionalized porphyrin [50] exhibits a linear stable response over a wide concentration range (1.0 × 10−5 to 1.0 × 10−1) with a slope of −57.4mV per decade, a detection limit of log[NO3−] = −5.72, and a selectivity coefficient for nitrate against perchlorate anion (logK pot NO3−,j = −1.81). Zinc-metallated 5, 10, 15, 20 –tetrakis (o-ferrocenylcarbonylaminophenylsubstituted)porphyrin atropisomers were prepared (see Structure 10) [51]. This zinc metalloporphyrins receptor strongly complex halide, nitrate and hydrogensulfate anions and stability-constant evaluations showed that these zinc metalloporphyrin receptor electrochemically sense anions via significant cathodic perturbations of the respective porphyrin oxidation and ferrocene redox couples. In the same field of porphyrins macrocycles, six polymer-modified electrodes were used to electrocatalyze the reduction of nitrate in two aqueous media [52]: 0.5 M HClO4 and 0.1 M NaClO4. These six polymers were prepared from the electropolymerization of tetrakis(xaminophenyl)porphyrin (where x = o; p) (see structure 11), and their Cu(II) and Ni(II) complexes in HCl solution on the surface of a conducting glass electrode. The results indicate that the catalytic response depends on the position of the amino group, the metal and the pH of the solution. In HClO4, the modified electrodes show two reduction processes.

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Structure 10. Structure taken from Scheme 1, reference 51 (Beer, P.; Drew, M.; Jagessar, R. J .Chem. Soc. 1997 881). Reproduced by permission of The Royal Society of Chemistry.

The voltammetric responses of the modified electrodes show a reduction process that takes place during the negative scan and another reduction process occurring during the return potential sweep. Indeed, there is a change in the number of reduction processes depending on the direction of the scan at low scan rate (5 mV s_1). In HClO4, the products of electrolysis are nitrite and nitrous oxide for Cu and Ni-ortho-polymers and nitrite for the others polymers. In NaClO4, only one reduction process takes place, independently of the direction of the scan, giving nitrite and nitrous oxide for the Ni and Cu ortho systems and only nitrite for the other systems. In both cases, HClO4 and NaClO4, the best polymers are poly Cu(II) and Ni(II)tetrakis (o-aminophenyl) porphyrin as electrocatalysts [52].

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H2N

N N

M

N

N NH2

NH2 Structure 11. M = Ni (II), Cu(II).

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CONCLUDING REMARKS Nitrate ion has relevance in the human activities, however their mismanaged may produce several harmful changes to the environment or human health. Therefore any analytical technique developed should be considered as an important contribution. In this scenario the use of coordination compounds appears as an alternative to develop and design modified electrodes with characteristics that allow the determination of this analyte. Because of the high overpotential required for the reduction of nitrate, the direct determination either by voltammetric or amperometric techniques has been poor exploited, however modified electrodes with macrocyclic compounds containing transition metals as potentiometric sensors could be the most feasible alternative. Indeed, in the course of this part of the chapter, it has been recognized that several families of macrocyclic compounds such as cyclams and porphyrins among others can be effectively used in the determination of nitrate achieving good values of parameters such as detection limit and selectivity.

I. 2. Nitrite Determination Nitrite ion formation is an important step in the nitrogen cycle. Nitrite is formed during the biodegradation of domestic or industrial nitrogenous wasted as well as some fertilizers. Nitrite ions are a precursor in the formation of nitrosamines, which have been shown to be carcinogenic [53]. Also, nitrite as contaminant, can cause many different type of noxious effect, such as microsomal enzyme inhibition and decrease in efficiency of nutritive diet [5456]. The nitrite levels in drinking water should be below 60 ngmL-1 [55]. Moreover, nitrites are routinely added to meat products as preservative against food poisoning microorganisms such as Clostridium Botulinum [57]. Thus, a number of different techniques for the

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quantitative determination of nitrite have been reported, which include electrochemistry [5860], chromatography [61-63], capillary electrophoresis [64, 65], spectrophotometry [66, 67] and spectrofluorimetry [68-78]. Moreover, electrochemical methods offer useful alternatives since they allow a faster and precise analysis. Particularly, electrochemical determination based on the oxidation of nitrite offers several advantages, namely no interference from nitrate ion and from molecular oxygen, which are usually the major limitation in cathodic determination of nitrite [79, 80]. Many important studies based on electrochemical nitrite oxidation on different kind of modified electrodes are presented below.

Structure 12. Molecular structure of the N,N´-bis(salicylidene)-4-methyl-1,2-phenylene diaminooxovanadium(IV). Structure taken from reference 81. Reprinted from Journal of Electroanalytical Chemistry, Volume 614(1-2), M. A.Kamyabi, F. Aghajanloo, Electrocatalytic oxidation and determination of nitrite on carbon paste electrode modified with oxovanadium(IV)-4-methyl salophen, page 158, Copyright (2008), with permission from Elsevier.

One interesting study was published by M.A. Kamyabi and D. Aghajanloo [81], and consist in the the electrocatalytic oxidation and detection of nitrite at carbon paste electrode chemically modified with N, N’-bis(salicylaldehyde)-4-methy-1,2 phenylenediimino oxovanadium (IV) ( VIVO(SB)-modified CPE) (see structure 12). The cyclic voltammetric response it shown in Figure (see Figure 1), were curves a and b correspond to bare carbon paste electrode, and curves c and d correspond to VIVO(SB)modified CPE, without and with nitrite (1mM) in 0.1 M phosphate buffer, pH4. It was found that in the bare electrode , anodic current due to the oxidation of nitrite is observed in b, but not cathodic peak it was found, indicating an irreversible charge transfer in that system. However when it was proved the modified electrode toward nitrite oxidation (curves c and d ), it was obtained a sharp anodic wave at lower positive potential for nitrite at the surface of VIVO(SB)-modified CPE, compared to unmodified electrode

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Figure 1. Cyclic voltammograms: at bare carbon paste electrode (a) in the absence, (b) in the presence of 1 mM nitrite; and at [oxovanadium-4-methyl salophen complex-modified carbon paste electrode] (c) in the absence and (d) in the presence of 1 mM nitrite; supporting electrolyte, 0.1M phosphate buffer, pH 4; scan rate 100 mV s-1. Figure taken from reference 81. Reprinted from Journal of Electroanalytical Chemistry, Volume 614(1-2), M. A.Kamyabi, F. Aghajanloo, Electrocatalytic oxidation and determination of nitrite on carbon paste electrode modified with oxovanadium(IV)-4-methyl salophen, page 160, Copyright (2008), with permission from Elsevier.

An increasing in the anodic peak current in conjunction with the sharpness of the peak, which is related to a decrease of the overpotential of the process, revealed that the modified electrode could act as an effective promoter to enhance the kinetics of the electrochemical process, it means that by using the VO(SB) as an electron mediator in the matrix of the modified electrode, the overpotential for the anodic oxidation of nitrite becomes considerably lower and the rate of the heterogeneous electron transfer in increased. The modified electrode used in this work was prepared by mixing unmodified composite (carbon paste electrode) with VO (salen) (25% w/w). Is interesting note that the amount of VO(SB) in the carbon paste has significant influence on the voltammetric response of the modified electrode. It was found that the oxidation current for 1mM nitrite increases gradually with the modifier, and at 25% (w/w) the oxidation current achieves a maximum and ten decreases with a further increase in the modifier (see Figure 2).

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Figure 2. Effect of modifier percentage on the peak current (circle) and peak potential (triangles) of cyclic voltammograms for 1 mM of nitrite in 0.1M phosphate buffer, pH 4. Figure taken from reference 81. Reprinted from Journal of Electroanalytical Chemistry, Volume 614(1-2), M. A.Kamyabi, F. Aghajanloo, Electrocatalytic oxidation and determination of nitrite on carbon paste electrode modified with oxovanadium(IV)-4-methyl salophen, page 160, Copyright (2008), with permission from Elsevier.

The authors attributed this behavior, due to a decreased in the graphite content in the paste and consequent reduction of the conductive electrode area. In the Figure it can see that the peak potential for the oxidation of nitrite is also affected by a change in the percentage of modifier, i.e. the peak potential is decreased between 0% and 25% modifier and then increases with increasing amount of modifier. Another interesting work for the nitrite oxidation corresponds to A.L. Sousa et al. [82]. It was presented an amperometric sensor for nitrite detection based on a glassy carbon electrode (GC) modified with copper tetrasulphinated phtalocyanine (CuTSPc) (see Structure 13) immobilized by polycationic poly-L-lysine film (PLL). The Figure compare the behaviour of bare GC (b), GC modified with PLL only (c), and GC modified with CuTSPc/PLL (d), in presence of nitrite 0.01 mmol L-1. Curve (a) show the voltammograms for GC modified with CuTSPc/PLL in absence of nitrite. As can be seen, the best response was obtained with the GC modified with CuTSPc/PLL (curve d). This electrode present higher peak current and lower potential for NO2- compared with all systems, showing a shift of 150 mV in the peak potential toward a less positive values. This behavior suggest that the CuTSPc complex is working as catalyst for NO2- oxidation on the sensor surface, and it associated with the low-charge resistance of the CuTSPc/PLL film. This electrochemical response toward nitrite electrooxidation depend of the pH when it was used 0.1 M of NO2- in buffer solution pHs 5, 6, 7 and 8 (see Figure 4), finding that the peak current reached a maximum value at pH 7.0 (Figure 4a), while the peak potential was independent of the pH (Figure 4b).

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O3S

SO3-2 N

N

N

Cu

N

N N

N

N -2

O3S

SO3-2

Structure 13.

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The voltametric response for nitrite electrooxidation mediated by a modified electrode with CuTSPc inmmobilized on PLL(CuTSPc/PLL), in phosphate buffer pH=7 is show in Figure3.

Figure 3. Cyclic voltammograms for CuTSPc/PLL modified electrode in the absence (a) and presence (d) of 0.25 mmol L−1 NO2−; for a bare GC electrode (b) and for an electrode with PLL only (c), both in presence of 0.25 mmol L−1 NO2−, in 0.1 mol L−1 phosphate buffer solution (pH 7.0). Scan rate: 0.05Vs−1. Inset: large-scale cyclic voltammograms for CuTSPc/PLL modified electrode in absence (a) and presence (d) of 0.25 mmol L−1 NO2−..Figure taken from reference 82. Reprinted from Talanta, Volume 75(2), A. L. Sousa, W. J.R. Santos, R. C.S. Luz, F. S. Damos, L.T. Kubota, A. A. Tanaka, S. M.C.N. Tanaka, Amperometric sensor for nitrite based on copper tetrasulphonated phthalocyanine immobilized with poly-l-lysine film, page 335, Copyright (2008), with permission from Elsevier.

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Figure 4. Influences of pH on the peak current (a) and on the peak potential (b) obtained by CV in 0.1 mol L−1 phosphate buffer containing 0.01 mmol L−1 NO2−. Scan rate: 0.015 Vs−1. Figure taken from reference 82. Reprinted from Talanta, Volume 75(2), A. L. Sousa, W. J.R. Santos, R. C.S. Luz, F. S. Damos, L.T. Kubota, A. A. Tanaka, S. M.C.N. Tanaka, Amperometric sensor for nitrite based on copper tetrasulphonated phthalocyanine immobilized with poly-l-lysine film, page 335, Copyright (2008), with permission from Elsevier.

This modified electrode was proved as amperometric sensor for nitrite. The sensor characteristics were verified by chronoamperometry, founding a best potential to be applied on the electrode. This applied potential was chosen based on measurements of the catalytic current intensities in the optimized conditions and the highest current was verified at an applied potential of 0.84 V versus SCE. The modified electrode presented a good repeatability for nitrite determinations, finding a linear response range from 0.12 up to 12.2 μmol L-1, with a good detection limit. This sensor was applied for nitrite determination in natural water samples, showing good responses compare with a spectrophotometric method. Also one interesting work is an amperometric nitrite sensor based on a polymeric nickel tetraaminophthalocyanine (p-NiTAPc) film coated glassy carbon (GC) electrode [83]. It was found that pH influenced determination of nitrite 0.1mM. The anodic peak current of nitrite increased with the decrease of the pH, founding that over the range of pH 1-3, higer sensitivity of voltammetric response of the p-NiTAPc film modified electrodes towards nitrite

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in H2SO4-Na2SO4 solution was obtained. The effect of thickness of the polymeric film on the anodic peak current of nitrite was evaluated. The film thickness it was controlled through the number of scans during the repetitive cyclic voltammetric polymerization of the p-NiTAPc film. It was found that the current response was small when the film was too thin due to the film possessing only few catalytically active sites. However when the film was too thick, a large background current which prevents the lowing of detection limit was observed. It was found that the optimum thickness of the p-NiTAPc film was obtained by scanning 6-10 cycles (see Figure 5).

Figure 5. Cyclic voltammogram of 0.1mM NiTAPc in DMF containing 0.1M TBAP at a glass carbon electrode. Scan rate: 100mVs−1. Figure taken from reference 83. Reprinted from Talanta, Volume 62(2), Z.-H.Wen, T.-F. Kang, Determination of nitrite using sensors based on nickel phthalocyanine polymer modified electrodes, page 353, Copyright (2004), with permission from Elsevier.

This modified electrode was proved as a nitrite sensor. It was found that the oxidation current was linearly related to nitrite concentration over the range 2.5 x 10-6 to 1.0 x 10-3 M, with a detection limit of 9 x 10-7 M. It was proved the precision of the method, where electrochemical experiment were repeatedly performed 12 times with the same p-NiTAPc modified electrode in the solution containing nitrite 1x 10-4M, showing a standard deviation of 0.022 and a relative standard of 2.8%, which revealed a good reproducibility of the modified electrode. Also it was found that the modified electrode was highly stable during 35 days only with a small decrease of current sensitivity (about 8% for nitrite 1 x 10-4 M). Another study of nitrite electrooxidation was the development of a higher sensitive amperometric sensor for nitrite using a glassy carbon electrode modified with alternated layer of iron (III) tetra-(N-methyl-4-pyridyl)-porphyrin (FeT4MPyP) and cobalt (II) tetrasulfonated phthalocyanine (CoTSPc) [84]. Figure 6 shows the electrocatalytic oxidation of nitrite on the modified electrode, where (a) correspond to layers of FeT4MPyP, (b) correspond to CoTSPc, and (c) to FeT4MPyP/CoTSPc.

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This sensor also was applied for nitrite determination in rivers water samples, showing good responses compare with a spectrophotometric method (see Table I). Table I. Comparison between a standard spectrophotometric method and the proposed method for nitrite determination

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Table taken from reference 83. Reprinted from Talanta, Volume 62(2), Z.-H.Wen, T.-F. Kang, Determination of nitrite using sensors based on nickel phthalocyanine polymer modified electrodes, page 355, Copyright (2004), with permission from Elsevier.

Figure 6. Cyclic voltammograms of the bare electrode (a) and modified electrodes with layers of: (b) FeT4MPyP, (c) CoTSPc and (d) FeT4MPyP/CoTSPc in presence of 0.5 mmol l−1 NO2−.Measurements carried out in 0.1 mol l−1 phosphate buffer. Scan rate: 0.05Vs−1. Figure taken from reference 84. Reprinted from Talanta, Volume 70(3), W.J.R. Santos , A.L. Sousa , R.C.S. Luz, F.S. Damos, L.T. Kubota , A.A. Tanaka, S.M.C.N. Tanaka, Amperometric sensor for nitrite using a glassy carbon electrode modified with alternating layers of iron(III) tetra-(N-methyl-4-pyridyl)-porphyrin and cobalt(II) tetrasulfonated phthalocyanine, page 590, Copyright (2006), with permission from Elsevier.

It can be seen, that the best response was obtained with the FeT4MPyP/CoTSPc modified electrode, showing higher peak current and lower oxidation potential for nitrite oxidation, compared with the modified electrode only with CoTSPc or FeT4MPyP. This electrode (FeT4MPyP/CoTSPc) was prepared by drop-casting transferring 7 μL of FeT4MPyP solutions of 0.1, 0.3, 0.5, 0.6 and 0.8mM concentrations, to the electrode surface, letting to dry at room temperature. Then, after 90 min, 7 μL of CoTSPc solutions of the same concentrations (0.1-0.8mM), was added on the electrode surface, letting to dry at room

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temperature during 90 min. Modified electrode with bilayers of complexes were prepared by this way. The modified electrode with CoTSPc only, presents oxidation potential, toward nitrite oxidation, close to the oxidation potential on the FeT4MPyP/CoTSPc modified electrode. However, the modified electrode with FeT4MPyP only, was not catalytic toward nitrite oxidation. The author suggests that the CoTSPc complex effectively catalyzes the nitrite oxidation, while the FeT4MPyP layers must be improving the assembly of the active sites of CoTSPc resulting in a better catalysis for nitrite oxidation reaction. It was found that the amount of FeT4MPyP and CoTSPc it was a control factor of great importance. This dependence was investigated by authors, preparing solutions containing different amounts of FeT4MPyP and CoTSPc (0.1-0.8 mM). Table II shows this study, indicating that the best response is obtained using 0.5mM.

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Table II. Influence of the FeT4MPyP and CoTSPc concentrations used on the preparation film on the peak current obtained with the sensor for 5 × 10− 4 mol L− NO2 in phosphate buffer solution (pH 7.0)

1

Table taken from reference 84. Reprinted from Talanta, Volume 70(3), W.J.R. Santos , A.L. Sousa , R.C.S. Luz, F.S. Damos, L.T. Kubota , A.A. Tanaka, S.M.C.N. Tanaka, Amperometric sensor for nitrite using a glassy carbon electrode modified with alternating layers of iron(III) tetra-(N-methyl4-pyridyl)-porphyrin and cobalt(II) tetrasulfonated phthalocyanine, page 590, Copyright (2006), with permission from Elsevier.

Figure 7 shows an analytical curve for the amperometric sensor, working under an applied potential of 0.85V versus SCE. Under optimization of experimental condition for amperometric yielded, it was found a detection limit and sensitivity for nitrite respectively, of 0.04μM and 0.37μLmol-1. This sensor showed good repeatability for the measurements and the relative standard deviation for 10 determination of 20μM of nitrite was 1.4%. This sensor also was applied in water samples, showing good result compared with spectrophotometric method described in literature [85] In other kind of modified electrodes, W.S. Cardoso and Y. Gushikem [86] studied the electrocatalytic oxidation of nitrite on a carbon paste electrode modified with Co(II) adsorbed on Si/O2SnO2/Phosphate prepared by sol-gel method. For the electrode preparation, the immobilization of the porphyrin was done by immersing 1 g of SiO2/SnO2/phosphate (previously prepared) in a 1mM solution of 5,10,15,20-tetrakis(1-methyl-4-pyridyl)-21-H, 23H-porphyrin tetra p-tosylate salt.

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Figure 7. Analytical curve for nitrite oxidation obtained in the optimized conditions in concentrations 0.2, 0.39, 0.58, 0.77, 0.95, 1.30, 1.62, 2.20, 2.61, 2.91, 3.19, 3.74, 4.25, 4.73, 5.19, 6.01, 6.75, 7.42 and 8.60 mol l−1 NO2− (inset: amperogram of the analytical curve). Figure taken from reference 84. Reprinted from Talanta, Volume 70(3), W.J.R. Santos , A.L. Sousa , R.C.S. Luz, F.S. Damos, L.T. Kubota , A.A. Tanaka, S.M.C.N. Tanaka, Amperometric sensor for nitrite using a glassy carbon electrode modified with alternating layers of iron(III) tetra-(N-methyl-4-pyridyl)-porphyrin and cobalt(II) tetrasulfonated phthalocyanine, page 592, Copyright (2006), with permission from Elsevier.

This mixture was shaken for 24 h and filtered washing the solid with deionized water, drying at room temperature. Furthermore, the adsorbed porphyrin species was metallated in situ by immersing 1.0 g of the material in 25 mL of a 1.0x10-2 M CoCl2 solution, and the suspension was shaken for 24 h. Then, Uv-Vis electronic absorption spectra for the solid powders containing unmetallated and metallated porphyrins were obtained by the diffuse reflectance spectroscopic technique. Evidence that the immobilized porphyrin is metallated was obtained from the changes in the Uv-Vis spectrum (see Figure 8). Non metallated porphyrin shows four bands, Q bands, at 502, 559, 588 and 643 nm (Figure 8a), and when it was metallated, they are reduced to two bands, one at 590nm and other at 570nm (Figure 8b), due to the increased of the local symmetry of the unmetallated porphyrin (D2h) to metallated porphyrin (D4h) [86]. This modified electrode was used toward nitrite oxidation. Figure 9 shows the electrocatalytic response toward nitrite oxidation. It was observed that an anodic oxidation peak current density to 0.72 V vs. SCE, and its intensity increased as the concentration of nitrite increased (see Figure 9). It was plotted the relationship between anodic oxidation peak current density against nitrite concentration, showing a linear correlation in the concentration range 1.0 x 10-5- 5.0 x 10-3M. For concentrations higher than 5.0x10-3 M it was observed electrode saturation (see Figure 10).

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Figure 8. Diffuse reflectance spectra for: (a) SiO2/SnO2/Phosphate/H2P; (b) SiO2/SnO2/Phosphate/CoP. Figure taken from reference 86. Reprinted from Journal of Electroanalytical Chemistry, Volume 583(2), W.S. Cardoso, Y. Gushikem, Electrocatalytic oxidation of nitrite on a carbon paste electrode modified with Co(II) porphyrin adsorbed on SiO2/SnO2/Phosphate prepared by the sol–gel method, page 302, Copyright (2005), with permission from Elsevier.

Figure 9. Cyclic voltammograms at a carbon paste electrode modified with SiO2/SnO2/Phosphate/CoP in 1.0 mol L-1 KCl in the absence of nitrite ions (1), in presence of 6.0 x 10-4 mol L-1 nitrite ions (2) and 1.0 x 10-3 mol L-1 nitrite ions (3). v = 0.02 Vs-1. Figure taken from reference 86. Reprinted from Journal of Electroanalytical Chemistry, Volume 583(2), W.S. Cardoso, Y. Gushikem, Electrocatalytic oxidation of nitrite on a carbon paste electrode modified with Co(II) porphyrin adsorbed on SiO2/SnO2/Phosphate prepared by the sol–gel method, page 303, Copyright (2005), with permission from Elsevier.

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Figure 10. Plot of jpc vs. mM of [NO-2] from cyclic voltammograms at a carbon paste electrode modified with SiO2/SnO2/Phosphate/CoP in 1.0 mol L-1 KCl for the concentration range between 1.0 x 10-5 and 7.0 x 10-3 mol L-1 of nitrite ions. v = 0.02 Vs-1, T = 298 K. Figure taken from reference 86. Reprinted from Journal of Electroanalytical Chemistry, Volume 583(2), W.S. Cardoso, Y. Gushikem, Electrocatalytic oxidation of nitrite on a carbon paste electrode modified with Co(II) porphyrin adsorbed on SiO2/SnO2/Phosphate prepared by the sol–gel method, page 303, Copyright (2005), with permission from Elsevier.

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In the field of phthalocyanines, B. Agboola and T. Nyokong studied the electroxidation of nitrite using Co(II), Fe(II) and Mn(II) tetrakis(benzylmercapto) and tetrakis (dodecylmercapto) phthalocyanines electrodeposited onto a gold electrode [87]. Structure 14 show the structures of metallophthalocyanine complexes used by authors, where CoTBMPc is called 1a; CoTDMPc 1b; FeTBMPc 2a; FeTDMPc 2b; MnTBMPc 3a and MnTDMPc 3b.

Structure 14. Molecular representative of the structures of metallophthacyanine complexes. Structure taken from reference 87. Reprinted from Analytica Chimica Acta, Volume 587(1), B. Agboola, T. Nyokong, Comparative electrooxidation of nitrite by electrodeposited Co(II), Fe(II) and Mn(III) tetrakis (benzylmercapto) and tetrakis (dodecylmercapto)phthalocyanines on gold electrodes, page 117, Copyright (2007), with permission from Elsevier.

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Figure 11 shows the comparative cyclic voltammograms for the catalytic activities of gold electrodes modified with these phthalocyanines, towards nitrite elctrooxidation. It was found that, based on the position of the catalytic peak potential, the increasing order for the complexes was: 3b < 1b