AP Chemistry 9.7: Electrolysis and Faraday’s Law

thermodynamically favored reaction

• converts chemical energy to electrical energy
• Ecell > 0 for favored direction
• ΔG < 0
• anode negative
• cathode positive

thermodynamically unfavored reaction is forced

• converts electrical energy into stored chemical energy
• requires an external power source
• anode positive
• cathode negative
• But again:
• oxidation is always at the anode
• reduction is always at the cathode

Electrolysis uses electrical energy to drive a redox reaction that would not proceed in the desired direction on its own.

• Common applications include:
• metal plating
• purification of metals
• production of substances from molten or aqueous ionic compounds

For AP Chemistry, the important part is usually the quantitative relationship between current, time, electrons, and chemical amount.

Current is:

• I = q/t
• Therefore:
• q = It
• where:
• I = current in amperes
• q = charge in coulombs
• t = time in seconds
• Remember:
• 1 ampere = 1 coulomb/second

One mole of electrons carries approximately:

• 96,485 C
• So:
• mol e⁻ = charge / F
• or:
• mol e⁻ = It/F

Then use the balanced half-reaction to convert electrons into the amount of substance produced or consumed.

For most quantitative electrolysis problems:

• current × time → charge → moles e⁻ → moles substance → grams

Suppose metal ion:

• M²⁺ + 2e⁻ → M
• If you determine that:
• 0.100 mol e⁻
• passed through the cell, then:
• 0.100 mol e⁻ × (1 mol M / 2 mol e⁻)
• = 0.0500 mol M

Then use molar mass to find grams.

If current is given in amperes:

• A = C/s
• so time should be converted to:
• seconds
• Example:
• 10.0 min × 60 s/min = 600 s

Do not automatically assume:

• 1 mol e⁻ = 1 mol metal
• Example:
• Al³⁺ + 3e⁻ → Al
• requires:
• 3 mol electrons per 1 mol Al
• while:
• Ag⁺ + e⁻ → Ag
• requires only:
• 1 mol electron per 1 mol Ag