Half‑equations in the hydrogen fuel cell

Energy changes • Chemical and fuel cells

Flashcards

Test your knowledge with interactive flashcards

Practical issue: What effect does fuel impurity have on electrode reactions?

Click to reveal answer

Fuel impurity causes side reactions or catalyst poisoning, reducing reaction efficiency and altering electrode performance.

Key concepts

What you'll likely be quizzed about

Definition of a half‑equation

A half‑equation shows either the oxidation or the reduction part of a redox reaction, listing species, electrons, and charges. Each half‑equation preserves mass and charge by balancing atoms and the net electrical charge using electrons and hydrogen or hydroxide ions as required. Half‑equations combine to form the full redox equation when the electrons lost in oxidation equal the electrons gained in reduction.

Anode reaction (acidic electrolyte)

Hydrogen oxidises at the anode and loses electrons, producing protons in an acidic electrolyte. The half‑equation in acidic conditions expresses hydrogen gas converting to protons and electrons: H2 -> 2H+ + 2e-. Loss of electrons causes current to flow through the external circuit away from the anode. Catalyst presence increases the rate of oxidation but does not alter the half‑equation.

Cathode reaction (acidic electrolyte)

Oxygen reduces at the cathode by accepting electrons and combining with protons to form water in acidic electrolyte. The balanced cathode half‑equation in acidic conditions is 1/2 O2 + 2H+ + 2e- -> H2O. Acceptance of electrons neutralises the charge of protons and results in water formation. Matching the two‑electron transfer at the cathode with the two‑electron loss at the anode yields an electrically balanced overall equation.

Half‑equations in alkaline electrolyte

Alkaline electrolytes conduct hydroxide ions instead of protons, changing the form of the half‑equations while conserving electron flow. Anode oxidation in alkaline conditions uses hydroxide ions: H2 + 2OH- -> 2H2O + 2e-. Cathode reduction produces hydroxide ions: 1/2 O2 + H2O + 2e- -> 2OH-. Combining these two half‑equations cancels electrons and intermediate water or hydroxide species to give the overall water‑forming reaction.

Balancing electrons and forming the overall equation

Electrons lost in the oxidation half‑equation must equal electrons gained in the reduction half‑equation. Multiplying half‑equations by suitable factors ensures electron balance before addition. In both acidic and alkaline forms, combining the two balanced half‑equations produces the overall reaction: 2H2 + O2 -> 2H2O. Conservation of mass and charge holds in the combined equation when electrons cancel.

Limiting factors that affect electrode reactions

Catalyst poisoning, fuel impurities, membrane hydration and temperature affect the rate and completeness of electrode reactions without changing the half‑equations. Catalyst poisoning decreases active sites and slows electron transfer. Impure hydrogen introduces side reactions that consume electrons or block the electrode. Low temperature reduces reaction rates while poor membrane hydration limits ionic conduction between electrodes.

Key notes

Important points to keep in mind

Half‑equations display either oxidation or reduction separately and must balance atoms and charge.

In acidic electrolyte, anode: H2 -> 2H+ + 2e-; cathode: 1/2 O2 + 2H+ + 2e- -> H2O.

In alkaline electrolyte, anode: H2 + 2OH- -> 2H2O + 2e-; cathode: 1/2 O2 + H2O + 2e- -> 2OH-.

Electrons lost at the anode equal electrons gained at the cathode before combining half‑equations.

Overall reaction always simplifies to 2H2 + O2 -> 2H2O regardless of electrolyte.

Catalyst, fuel purity, temperature and membrane condition affect reaction rates but do not change half‑equations.

Balance oxygen and hydrogen atoms using H2O, H+ or OH- depending on the electrolyte.

Membrane type determines whether H+ or OH- transports between electrodes.