1.3.5—Fuel cells

A fuel cell converts chemical energy from a spontaneous redox reaction directly into electrical energy. Oxidation occurs at the anode and reduction at the cathode.

Deduce each half-equation from the fuel-cell reactants and electrolyte context, balance atoms and charge, then add the half-equations to obtain the overall reaction. Hydrogen and methanol cells require different oxidation half-equations.

For a hydrogen fuel cell, oxidation of H₂ supplies electrons at the anode and O₂ gains electrons at the cathode; the half-equations must match the acidic or alkaline electrolyte before they are added to 2H₂ + O₂ → 2H₂O. Direct electrical conversion does not remove the need to evaluate fuel production and storage.

A fuel cell operates while fuel and oxidant are supplied continuously from outside; a conventional battery stores a finite set of reactants internally. Point-of-use water from a hydrogen cell is not a complete environmental assessment—fuel manufacture, transport, storage and electricity source remain inside a lifecycle comparison.

Worked equations — acidic hydrogen cell: anode $\ce{H2(g) -> 2H+(aq) + 2e-}$; cathode $\ce{O2(g) + 4H+(aq) + 4e- -> 2H2O(l)}$. Double the anode equation before adding, giving $\ce{2H2 + O2 -> 2H2O}$. Direct-methanol cell: anode $\ce{CH3OH(aq) + H2O(l) -> CO2(g) + 6H+(aq) + 6e-}$; cathode $\ce{3/2O2(g) + 6H+(aq) + 6e- -> 3H2O(l)}$. Adding and cancelling gives $\ce{CH3OH + 3/2O2 -> CO2 + 2H2O}$. Match each half-equation to the stated electrolyte; proton-exchange-membrane construction details are not assessed.