Electrolytic Cells

Core Concept

An electrolytic cell uses an external source of electrical energy to drive a nonspontaneous redox reaction — the reverse of a galvanic cell. The external supply forces electrons to move in the opposite direction, so the electrode signs are reversed (anode positive, cathode negative in electrolysis), but the chemistry is unchanged: oxidation still occurs at the anode and reduction at the cathode. Electrolysis decomposes compounds (molten salts to metals, water to H₂ and O₂) and is the industrial basis of aluminum, chlorine, and electroplating.

Key Ideas

• Nonspontaneous, powered. The cell reaction has ΔG > 0 (E°cell < 0); the external supply overcomes it.
• Signs flip, roles don't. In an electrolytic cell the anode is the positive electrode and the cathode negative, but oxidation is still at the anode and reduction at the cathode.
• Electrons pumped backward. The external source pushes electrons onto the cathode (reduction) and pulls them from the anode (oxidation).
• Molten salts give the elements. Electrolysis of molten NaCl yields Na metal and Cl₂ gas.
• Aqueous solutions compete with water. In water, H₂O can be reduced (to H₂ + OH⁻) or oxidized (to O₂ + H⁺) instead of, or in competition with, the solute ions.

Equations and Variables

How It Works

1. An external power supply is connected to two electrodes immersed in an electrolyte (molten salt or solution).
2. The supply's negative terminal feeds electrons to the cathode, where cations are reduced.
3. The supply's positive terminal draws electrons from the anode, where anions (or the solvent) are oxidized.
4. The applied voltage must exceed the spontaneous back-voltage (plus any overpotential) for current to flow.
5. The result is a net nonspontaneous decomposition driven by electrical work.

Worked Example

Predict the products of electrolysis of molten NaCl, and label the electrodes.

Molten NaCl contains Na⁺ and Cl⁻ (no water to compete).

• Cathode (reduction): Na⁺ + e⁻ → Na(l) — sodium metal forms.
• Anode (oxidation): 2 Cl⁻ → Cl₂(g) + 2 e⁻ — chlorine gas forms.
• Net: 2 NaCl(l) → 2 Na(l) + Cl₂(g).

For aqueous NaCl, water competes: at the cathode, water is more easily reduced than Na⁺, so 2 H₂O + 2 e⁻ → H₂ + 2 OH⁻ (hydrogen, not sodium); at the anode, 2 Cl⁻ → Cl₂ + 2 e⁻ (chlorine). So aqueous NaCl electrolysis gives H₂ and Cl₂ (with NaOH left in solution), not sodium metal.

How it works

1. An external power supply is connected to two electrodes immersed in an electrolyte (molten salt or solution).
2. The supply's negative terminal feeds electrons to the cathode, where cations are reduced.
3. The supply's positive terminal draws electrons from the anode, where anions (or the solvent) are oxidized.
4. The applied voltage must exceed the spontaneous back-voltage (plus any overpotential) for current to flow.
5. The result is a net nonspontaneous decomposition driven by electrical work.

Common confusions

• "In an electrolytic cell the anode does reduction." — Wrong. Anode = oxidation and cathode = reduction in every cell, galvanic or electrolytic.
• "The anode is always negative." — Wrong. In electrolysis the anode is positive (connected to the supply's + terminal); only in a galvanic cell is the anode negative.
• "Electrolyzing aqueous NaCl gives sodium metal." — Wrong. Water is reduced to H₂ instead; sodium metal requires molten NaCl (no water).
• "No voltage is needed if the reaction is spontaneous." — Then it wouldn't be electrolysis; electrolysis is for nonspontaneous reactions, so external voltage is required.
• "The cathode and anode trade which reaction they host." — Wrong. They trade electrical sign, not their chemical roles.

Quick review

• Electrolytic cell: nonspontaneous reaction driven by external electricity.
• Anode = oxidation (+), cathode = reduction (−) — signs reversed vs. galvanic.
• Molten salt → elements; aqueous → water may compete.
• Aqueous NaCl → H₂ + Cl₂ (not Na).
• Applied voltage must exceed the back-voltage.