How Electrolysis Works

Part of Electrochemistry

The fundamental principles of electrolysis — how electrical current drives chemical reactions at electrode interfaces.

Why This Matters

Electrolysis is one of the most powerful tools in applied chemistry: it uses electrical energy to drive chemical reactions that would not occur spontaneously at normal temperatures and pressures. Without electrolysis, there is no aluminum production, no sodium hydroxide, no chlorine for water treatment, no electroplating, no high-purity metals for electronics, and no practical hydrogen production for fuel.

Understanding how electrolysis works at a fundamental level — what ions do, what the electrodes do, how potential drives reactions — lets you apply the principle to novel problems. You are not limited to following recipes; you can design new processes when the situation demands.

The Basic Setup

An electrolysis cell has five essential components:

1. Power supply (DC, not AC — direction matters)
2. Anode — the positive electrode
3. Cathode — the negative electrode
4. Electrolyte — an ionic conductor between the electrodes
5. External circuit — the wire connecting anode and cathode through the power supply

Without any one of these, the circuit is open and nothing happens.

Ionic Conduction in the Electrolyte

Metals conduct electricity through free electrons. Electrolytes conduct through the movement of ions — charged atoms or molecules that carry charge through the liquid.

When an ionic compound (like copper sulfate, CuSO₄) dissolves in water, it dissociates: CuSO₄ → Cu²⁺ + SO₄²⁻

In the electric field between anode (+) and cathode (−):

• Cations (positively charged ions, Cu²⁺) migrate toward the cathode (negative electrode)
• Anions (negatively charged ions, SO₄²⁻) migrate toward the anode (positive electrode)

This movement of ions is the electric current in the electrolyte. The total current is the same in the wire and in the electrolyte — charges in, charges out, circuit complete.

Electrode Reactions

When ions reach an electrode, electrochemical reactions occur. The type of reaction depends on whether the electrode is anode or cathode:

At the cathode (reduction): The electrode supplies electrons to ions or molecules arriving at its surface. Species gain electrons — reduction.

Cu²⁺ + 2e⁻ → Cu° (copper deposits) 2 H⁺ + 2e⁻ → H₂ (hydrogen gas evolves)

At the anode (oxidation): The electrode withdraws electrons from species at its surface. Species lose electrons — oxidation.

Cu° → Cu²⁺ + 2e⁻ (copper anode dissolves) 2 Cl⁻ → Cl₂ + 2e⁻ (chlorine gas evolves) 2 H₂O → O₂ + 4H⁺ + 4e⁻ (oxygen gas evolves)

The specific reactions that occur depend on what species are present and the electrode potential.

Electrode Potential and the Electrochemical Series

Different species require different electrical potentials to be reduced or oxidized. The standard electrode potential (E°, measured in volts relative to the hydrogen electrode) quantifies this:

Half-Reaction	E° (V)
Li⁺ + e⁻ → Li	−3.04
Al³⁺ + 3e⁻ → Al	−1.66
Zn²⁺ + 2e⁻ → Zn	−0.76
Fe²⁺ + 2e⁻ → Fe	−0.44
Ni²⁺ + 2e⁻ → Ni	−0.25
2H⁺ + 2e⁻ → H₂	0.00 (reference)
Cu²⁺ + 2e⁻ → Cu	+0.34
Ag⁺ + e⁻ → Ag	+0.80
Au³⁺ + 3e⁻ → Au	+1.50

More positive E° = more easily reduced (deposits preferentially at cathode). More negative E° = requires more driving voltage to deposit.

Practical implication: In a solution containing Cu²⁺ and Zn²⁺, copper will deposit preferentially (more positive E°). To deposit zinc, you need to go to a more negative cathode potential, where zinc deposition competes favorably — or use a special bath chemistry that modifies the effective potential.

Minimum Voltage Required

To drive an electrolytic reaction, the power supply must provide at least the thermodynamic minimum voltage (the cell decomposition voltage) plus additional overpotential at each electrode:

V_cell = E_cathode − E_anode + η_cathode + η_anode + IR_ohmic

Where:

• E_cathode, E_anode = equilibrium potentials for the half-reactions
• η = overpotentials (extra voltage needed to drive reactions at practical rates)
• IR_ohmic = voltage drop in electrolyte (current × resistance)

For water electrolysis:

• Thermodynamic minimum: 1.23 V
• Typical actual cell voltage: 1.7–2.0 V (due to overpotentials and resistance)

For copper electrorefining:

• Near-zero thermodynamic driving force (same reaction at anode and cathode)
• Actual cell voltage: 0.2–0.35 V (almost all is ohmic drop in electrolyte)

DC vs. AC

Electrolysis requires DC (direct current). With AC (alternating current):

• Deposits form during one half-cycle and dissolve during the other half-cycle
• Net deposit is zero
• The process oscillates but produces nothing useful

The rectifier that converts AC to DC is thus a critical component of any industrial electrolytic facility. In a rebuilding context, generating DC directly (battery, DC generator) or building a rectifier (selenium stack, diode bridge from vacuum tubes or solid-state diodes) is necessary before electrolytic processes can be operated.

Faraday’s Laws in Brief

The amount of substance produced at an electrode is exactly proportional to the charge passed:

m = (M × I × t) / (F × z)

Where M is molar mass, I is current (amps), t is time (seconds), F is Faraday’s constant (96,485 C/mol), and z is the number of electrons per ion.

This law is exact — no exceptions in any electrolytic process. It is the foundation of all quantitative electrochemical design.

What Electrolysis Cannot Do (Easily)

Some reactions require such extreme potentials that aqueous electrolysis cannot reach them — the water itself electrolyes at lower potential. To deposit alkali metals (sodium, potassium), alkaline earth metals (calcium, magnesium), or aluminum, you must either use molten salt electrolytes (no water to interfere) or special ionic liquid systems. This is why aluminum production requires the Hall-Héroult molten salt process rather than simple aqueous electrolysis.