The simplest picture of a fuel cell is a magic box that gives you electricity, plus a bit of heat and water, when you feed it with hydrogen gas and oxygen. Ordinary air can supply enough oxygen, so all that is really needed is a source of pure hydrogen gas.
You may be curious enough to want to know what exactly goes on inside such a box. If so then read this detailed account provided by Dr Rob Potter of the Johnson Matthey Technology Centre:
If you pass a direct current through acidified or alkali water, you get hydrogen gas released at the negative (reducing) electrode and oxygen gas released at the positive (oxidising) electrode. With the right type of electrodes, if you turn the power supply off and, instead, connect the two electrodes to e.g. a small light-bulb, you can show that the process is somewhat reversible. The hydrogen and oxygen bubbles now go back into solution as water and you get power back out. This is what Sir William Grove discovered back in the 1830’s using platinum electrodes. This is the principle behind the fuel cell.
But what is actually going on at a molecular level, and how does nature allow us to get electrical energy from chemicals rather than just by burning them in ordinary combustion?
The hydrogen + oxygen reaction
H2 + ½ O2 = H2O …. (1)
is a ‘downhill’ reaction, which means that energy will be liberated during the formation of the product water. Chemists can easily calculate how much free energy should be released – around 237 kJ per mole of water. Both hydrogen and oxygen exist as di-atomic molecules – they prefer to be in pairs hence our formulae in reaction (1) are written as H2 and O2. In the gas-phase reaction, hydrogen molecules collide with oxygen molecules, and both hydrogen and oxygen molecules split apart and re-arrange into molecules of H2O. Energetically, this looks difficult to do as the strength of the dihydrogen bond is around 436 kJ/mol and that of dioxygen some 498 kJ/mol. Nonetheless, the gas phase reaction can occur very readily, often with explosive force.
In a fuel cell, the same reaction as the gas phase one is taking place, only this time the hydrogen and oxygen gas molecules are kept separate by a polymer membrane and are only allowed to collide with electrode surfaces. The same diatomic bonds need to be broken, but the intermediates that are formed at the electrode surfaces are different from the gas phase:
H2 = 2H+ + 2e- ….. (2)
½ O2 + 2e- = 2O2- …..(3)
At the oxidising electrode (anode), the dihydrogen gas is ionised to form protons - reaction (2). This is for an acid fuel cell, alkali fuel cells work in a similar way but with different solution intermediates. The protons are dissolved in water and move (by migration and diffusion) through the special polymer membrane towards the reducing electrode (cathode). Here the protons collide with oxygen species made by reaction (3) (note that the oxygen reaction is more complicated than reaction (3) implies, and the mechanism is still not properly understood). Electrons are injected into the oxygen intermediates via the special electrode (catalytic) coating and water molecules are formed. The 273 kJ/mol of energy associated with this reaction now appears as useful electrical work done by the electrons that have gone into the circuit at the oxidising electrode and re-appeared at the reducing electrode. Little heat is evolved in contrast to the gas phase reaction where almost all the energy is released thermally.
This is, as you will realise, a very simplified picture of the way the (acid polymer) fuel cell works at a molecular level. The two most important concepts to appreciate are that:
1) Nature does allow us to extract energy from a chemical reaction in different ways even though the end product is the same, and;
2) In a fuel cell, the reactant gases are not allowed to collide with each other, they collide with separate electrode surfaces and so the collision dynamics and types of key intermediates involved are often very different from those in the gas-phase reaction.
Johnson Matthey Technology Centre