Fuel Cell Systems Explained

Andrew L. Dicks, David A. J. Rand

• Éditeur: Wiley
• Année: 2018
• ISBN:
  9781118706961
• Édition: 3
• Sujet: Technologie et ingénierie
• Sous-sujet: Ressources d'alimentation

Since publication of the first edition of Fuel Cell Systems Explained, three compelling drivers have supported the continuing development of fuel cell technology. These are: the need to maintain energy security in an energy-hungry world, the desire to move towards zero-emission vehicles and power plants, and the mitigation of climate change by lowering of CO 2 emissions. New fuel cell materials, enhanced stack performance and increased lifetimes are leading to the emergence of the first truly commercial systems in applications that range from fork-lift trucks to power sources for mobile phone towers. Leading vehicle manufacturers have embraced the use of electric drive-trains and now see hydrogen fuel cells complementing advanced battery technology in zero-emission vehicles. After many decades of laboratory development, a global but fragile fuel cell industry is bringing the first commercial products to market.

This thoroughly revised edition includes several new sections devoted to, for example, fuel cell characterisation, improved materials for low-temperature hydrogen and liquid-fuelled systems, and real-world technology implementation.

Assuming no prior knowledge of fuel cell technology, the third edition comprehensively brings together all of the key topics encompassed in this diverse field. Practitioners, researchers and students in electrical, power, chemical and automotive engineering will continue to benefit from this essential guide to the principles, design and implementation of fuel cell systems.

1
Introducing Fuel Cells

1.1 Historical Perspective

This book is an introduction to fuel‐cell systems; it aims to provide an understanding of the technology — what it is, how it works and what are its applications. Essentially, a fuel cell can be defined as a device that produces electrical power directly from a fuel via an electrochemical process. In some respects, this operation is similar to that of a conventional battery except that the reactants are stored outside the cell. Therefore, the performance of the device is limited only by the availability of the fuel and oxidant supply and not by the cell design. For this reason, fuel cells are rated by their power output (kW) rather than by their capacity (kWh).

Before addressing the technology in depth, it is necessary to understand that by virtue of being electrochemical, fuel cells have both chemical and electrical characteristics. Accordingly, their development has been inextricably linked with the development of electrochemistry as a distinct branch of physical chemistry.

At the beginning of the 19th century, it was recognized that an ‘electrochemical cell’ (nowadays, commonly called a ‘battery’) could be made by placing two dissimilar metals in an aqueous salt solution. This discovery was made by Alessandro Volta, the professor of experimental physics at Pavia University, who constructed a pile of alternating discs of copper (or silver or brass) and zinc (or tin) that were separated by pasteboard discs (or ‘any other spongy matter’) soaked in brine. When the top and bottom of the pile were connected by a wire, the assembly delivered, for the first time in history, a more or less steady flow of electricity. Volta introduced the terms ‘electric current’ and ‘electromotive force’, the latter to denote the physical phenomenon that causes the current to flow. In due course, he conveyed his findings in a letter dated 20 March 1800 to Joseph Banks, the then president of the Royal Society. Known as the ‘Volta (or Voltaic) pile’, this was the first ‘primary’ (or non‐rechargeable) power source, as opposed to a ‘secondary’ (or rechargeable) power source.

Sir Humphry Davy, who was working at the Royal Institution in London, soon recognized that the Volta pile produces electricity via chemical reactions at the metal–solution interfaces — hydrogen is evolved on the ‘positive’ copper disc, and zinc is consumed at the ‘negative’ disc. Indeed, this recognition of the relationship between chemical and electrical effects prompted Davy to coin the word ‘electrochemical’, from which sprang the science of ‘electrochemistry’. He gave warning that Volta’s work was ‘an alarm bell to experimenters all over Europe’. His prediction was soon to be verified.

Volta had sent his letter to the Royal Society in two parts because he anticipated problems with its delivery given that correspondence from Italy had to pass through France, which was then at war with Britain. While waiting for the second part to arrive, Joseph Banks had shown the first few pages to Anthony Carlisle (a fashionable London surgeon) who, in turn, with the assistance William Nicholson (a competent amateur scientist) assembled on 30 April 1800 the first pile to be constructed in England. Almost immediately, on 2 May 1800, the two investigators found that the current from their device when passed through a dilute salt solution via two platinum wires was capable of decomposing water into its constituents — hydrogen at one wire and oxygen at the other. Details of the discovery were published in Nicholson’s own journal in July of the same year. Thus, the new technique of ‘molecular splitting’ — to be coined ‘electrolysis’ by Michael Faraday much later in 1834 and derived from the Greek ‘lysis’ = separation — was demonstrated before Volta’s own account of the pile was made public in September 1800. A schematic representation of the process is shown in Figure 1.1a.

Figure 1.1 Terminology employed in operation of (a) electrolysis cells and (b) fuel cells.

It was left to Michael Faraday, Davy’s brilliant student, to identify the mechanisms of the processes that take place within ‘electrolytic’ cells and to give them a quantitative basis. In addition, he was also the guiding force behind the nomenclature that is still in use today. First, Faraday with the assistance of Whitlock Nicholl (his personal physician and accomplished linguist) devised the name ‘electrode’ to describe a solid substance at which an electrochemical reaction occurs and ‘electrolyte’ to describe the chemical compound that provides an electrically conductive medium between electrodes. (Note that in the case of dissolved materials, it is fundamentally incorrect to refer to the ‘electrolyte solution’ as the ‘electrolyte’; nevertheless, the latter terminology has become common practice.) To distinguish between the electrode by which conventional current (i.e., the reverse flow of electrons) enters an electrolytic cell and the electrode by which it leaves, Faraday sought the assistance of the polymath William Whewell, the Master of Trinity College at the University of Cambridge. In a letter dated 24 April 1834, he asked Whewell:

‘Can you help me out to two good names not depending upon the idea of a current in one direction only or upon positive or negative?’

In other words, he wanted terms that would be unaffected by any later change in the convention adopted for the direction of current. Eventually, they settled on calling the positive electrode an ‘anode’ and the negative electrode a ‘cathode’, which were coined from Greek ‘ano‐dos’ (‘upwards’–‘a way’) to represent the path of electrons from the positive electrode to the negative and ‘katho‐dos’ (‘downwards’–‘a way’) to represent the counter direction. For an electrolytic cell, then, the anode is where the current enters the electrolyte and the cathode is where the current leaves the electrolyte. Thus the positive electrode sustains an oxidation (or ‘anodic’) reaction with the liberation of electrons, while a reduction (or ‘cathodic’) reaction takes place at the negative electrode with the uptake of electrons.

With use of the Greek neutral present participle ‘ion’ — ‘a moving thing’ — to describe the migrating particles in electrolysis, two further terms were obtained, namely, ‘anion’, i.e., the negatively charged species that goes to the anode against the current (or with the flow of negative charge), and ‘cation’, i.e., the positively charged species that goes to the cathode with the current (or against the flow of negative charge). The operation of an electrolysis cell is shown in Figure 1.1a. It should be noted that the anode–cathode terminology for an ‘electrolytic cell’ applies to a ‘battery under charge’ (secondary system).

A fuel cell operates in the reverse manner to an electrolysis cell, i.e., it is a ‘galvanic’ cell that spontaneously produces a voltage (similar to a ‘battery under discharge’). The anode of the electrolysis cell now becomes the cathode and the cathode becomes the anode; see Figure 1.1b. Nevertheless, the directions of the migration of anions and cations with respect to current flow are unchanged such that the positive electrode remains a positive electrode and the negative electrode remains a negative electrode. T...