Oxygen Reduction Catalysts for Fuel Cells

Christopher Barile
October 23, 2010

Submitted as coursework for Physics 240, Stanford University, Fall 2010

Fig. 1: Cell voltage versus the current density in a typical hydrogen fuel cell. [4]

Fuel cells have received increasing attention over the last decade for their potential to be a key component of a future green energy system. Since the only byproducts of a hydrogen fuel cell are water and waste heat, fuel cells provide a route of generating energy devoid of greenhouse gas emissions as long as the hydrogen consumed is produced in a clean manner. Many people are calling for a solar-hydrogen revolution, in which photovoltaic cells provide the electricity needed to electrolyze water to synthesize hydrogen [1]. In this energy utopia, photovoltaic cells would replace coal power plants while fuel cells would replace internal combustion engines.

Despite the recent flurry of interest, fuel cell technology has a long history. In 1839, William Grove constructed the first fuel cell using hydrogen and oxygen fuels and sulfuric acid as an electrolyte. [2] The first important commercial application of fuel cells came in the 1960s when they were used onboard the spacecrafts of Project Gemini. [3]

A fuel cell is a device that converts the chemical energy of a fuel and an oxidant into electricity. The two electrodes of the cell are separated by an electrolyte or membrane and contain the catalysts required to drive the necessary redox chemistry. The oxidant is almost always oxygen, whereas the fuel can be methanol, ethanol, or formic acid, although hydrogen is used mostly commonly. Perhaps the most attractive aspect of fuel cells is that their thermodynamic efficiencies are not limited by the Carnot cycle as is the case with combustion engines. Instead, fuel cell efficiency is directly related to the overpotential, the difference between the experimental voltage and thermodynamically prescribed voltage, of the half reactions at each electrode.

Whereas burning hydrogen in an internal combustion engine gives efficiencies ranging from 10-20%, the thermodynamic efficiency of a hydrogen fuel cell can be greater than 90%. [4] However, in practice such high efficiencies are never achieved. State-of-the-art fuel cells for vehicles have efficiencies near 50%, but this number does not reflect losses that come from making the hydrogen nor storing it.

The most common fuel cell catalyst at both electrodes is platinum. This noble metal along with its congeners nickel and palladium facilely oxidize hydrogen to hydrogen ions at the anode. The overpotential for this half reaction is typically only 50 mV. In contrast, the oxygen reduction reaction that occurs at the cathode usually has an overpotential of 500-600 mV with a platinum-based catalyst. This large overpotential at the cathode results in a substantial loss in efficiency, and thus the development of superior oxygen reduction catalysts must be addressed before even considering issues of hydrogen storage and production.

The catalytic four electron reduction of oxygen to water is kinetically slow because of the very strong double bond of molecular oxygen that must first be broken. Fig. 1 shows that for a typical hydrogen fuel cell, the cell voltage decreases with increased current output. Due to the high overpotential of the oxygen reduction reaction, substantial current densities cannot be achieved at voltages even close to the thermodynamic value of 1.2 V. At low current densities, there is a rapid decrease in cell voltage with increased current due to the slow kinetics of the cathodic catalysis. At current densities between about 0.2-1.0 A/cm2, the voltage declines linearly with increasing current due to normal Ohmic resistance. At higher current densities, the voltage falls off rapidly because the gasses cannot diffuse to the electrodes as fast as the reaction on the catalyst surface occurs.

For commercial fuel cells, a current density of about 1.5 A/cm2 of electrode material is normally desirable. High catalysts loadings are needed to achieve current densities of this magnitude and with platinum, cost becomes a significant barrier. In order to minimize the amount of catalyzed utilized, platinum nanoparticles physisorbed onto porous carbon surfaces have been developed in order to increase the catalytic site surface area. It might seem that thus a fuel cell with a cathode consisting of the smallest possible nanoparticles would be the most efficient. However, thorough research has demonstrated that the catalytic activity per surface area of platinum nanoparticles, or that of any catalytically active metallic nanoparticles for that matter, reaches a definite maximum. The catalytic activity of platinum nanoparticles reaches a maximum at a particle diameter of 2 nm. [5]

There have been many studies aimed at elucidating the mechanistic origin of the decreased per surface catalytic activity observed at smaller particle diameters, and the issue remains a matter of debate. Most researchers have reached the consensus that the lower catalytic activity comes as a consequence of a larger percentage or platinum atoms being on edge sites [6]. There is evidence that suggests that oxygen reduction preferentially occurs at sites where multiple platinum atoms can participate in catalysis. [7] Some material scientists, however, contend that the decreased catalytic activity obtained from platinum nanoparticles of diameters less than 2 nm has to do with limitations by mass transport due to decreased interparticle distance in the catalyst structure. [8]

Since the surface area of a platinum catalyst cannot be increased endlessly without diminishing returns, other strategies must be employed in order to decrease platinum loadings in the cathodes of fuel cells. It has been found that numerous platinum alloys exhibit comparable or even higher catalytic oxygen reduction activity than platinum. Nickel and cobalt alloys of platinum have shown the most promise. [9-11]

Another method that has been used to decrease the amount of platinum required in the electrode is to construct a shell of platinum around some other metal. [12-14] The development of nanoparticles consisting of cores and shells is at the frontier of oxygen reduction catalyst research. Nanoparticles with shells of platinum and cores of ruthenium, palladium, or even platinum alloys have been synthesized. However, these core metals are just as expensive as the platinum shells; nanoparticles with cores of cheaper metals have proved elusive [5].

Despite the extensive research conducted in the last decade focusing on designing platinum nanoparticles for the catalytic reduction of oxygen to water, the overpotential for this reaction with even the best catalysts remains prohibitively high at the current densities required for efficient, reasonably priced, commercial fuel cells. [4,5] The main obstruction to further advances in this field is that the details of the mechanism of the oxygen reduction reaction catalyzed by a metallic surface are exceedingly difficult to ascertain. [15-17] Numerous other nonprecious metal oxygen reduction catalysts have been designed including carbon nanotubes, porphyrins, and phthalocyanines, but organic-based catalysts such as these have thus far suffered from even higher overpotentials and instability. [4,18,19] The lack of an efficient oxygen reduction catalyst remains the paramount obstacle to the widespread application of fuel cells.

© 2010 Christopher Barile. The author grants permission to copy, distribute and display this work in unaltered form, with attribution to the author, for noncommercial purposes only. All other rights, including commercial rights, are reserved to the author.


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