A fuel cell is a galvanic cell supplied continuously with reactants. Oxidation occurs at its negative anode, reduction at its positive cathode, and electrons travel through the external circuit from anode to cathode. Ions move through the electrolyte to maintain charge balance. In an acidic hydrogen cell: H2→2H++2e− at the anode and O2+4H++4e−→2H2O at the cathode.
The overall reaction is 2H2+O2→2H2O. Continuous supply distinguishes fuel cells from sealed cells with fixed reactant inventories; flow cells can store reactants externally. Photovoltaic cells convert light directly to electrical energy. Hydrogen is an energy carrier whose production matters: water electrolysis uses electricity, while fossil-based production can release CO2. Storage, infrastructure, efficiency and electricity source affect comparisons.
Example 1
Combine the acidic hydrogen-cell half-equations.
Double H2→2H++2e− to give 2H2→4H++4e−. Add O2+4H++4e−→2H2O and cancel H+ and e−. Overall: 2H2+O2→2H2O.
Example 2
How many moles electrons pass through the circuit when 0.20 mol H2 is oxidised?
H2→2H++2e− gives two moles electrons per mole H2. Thus 0.20×2=0.40 mol electrons.
Watch out:The cell may produce water locally, but hydrogen production and energy supply can cause emissions elsewhere.
Fuel cells, hydrogen and direct electricity practice | Year 12 Chemistry | Ascent Arcade