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Table 2-2 Fuel Cell Reactions and the Corresponding Nernst Equations Cell Reactions* H2 +1⁄2O2 → H2O CO+1⁄2O2 →CO2 CH4 + 2O2 → 2H2O+ CO2 (a) - anode P (c) - cathode R E - equilibrium potential T F - Faraday's constant Nernst Equation E=E°+(RT/2F)ln[P /P ]+(RT/2F)ln[P1⁄2] H2 H2 O O2 E=E°+(RT/2F)ln[P /P ]+(RT/2F)ln[P1⁄2] CO CO2 O2 E=E°+(RT/8F)ln[P /P2 P ]+(RT/8F)ln[P2 ] CH4 H2O CO2 O2 - gas pressure - universal gas constant - temperature (absolute) H2+1⁄2O2+CO2(c) → H2O+ CO2(a) E = E° + (RT/2F) ln[PH2 /PH2O(PCO2 )(a)] + (RT/2F)ln[P1⁄2 (P ) ] O2 CO2 (c) * The cell reactions are obtained from the anode and cathode reactions listed in Table 2-1. The ideal standard potential (Eo) at 298K for a fuel cell in which H2 and O2 react is 1.229 volts with liquid water product, or 1.18 volts with gaseous water product. This value is shown in numerous chemistry texts (2) as the oxidation potential of H2. The potential is the change in Gibbs free energy resulting from the reaction between hydrogen and oxygen. The difference between 1.229 volts and 1.18 volts represents the Gibbs free energy change of vaporization of water at standard conditions. Figure 2-1 shows the relation of E to cell temperature. Because the figure shows the potential of higher temperature cells, the ideal potential corresponds to a reaction where the water product is in a gaseous state (i.e., Eo is 1.18 volts). 1.2 1.1 1.0 Figure 2-1 H2/O2 Fuel Cell Ideal Potential as a Function of Temperature 300 400 500 600 700 800 900 1000 1100 Temperature (K) 2-5 H 2 + 12 O 2 H 2 O ( g ) Reversible potential (V)PDF Image | Fuel Cell Handbook (Seventh Edition)
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