IMAT Chemistry · Redox Reactions

Electrochemical Cells and Electrolysis

A galvanic cell lets a spontaneous redox reaction push electrons through a wire; an electrolytic cell uses a power supply to drive a redox reaction that would not happen by itself.

All Redox Reactions notes

Why this matters

One older question, from 2014, asked for the oxidation number changes in a car battery as it discharges. Cells and electrolysis are on the syllabus at outline level; expect facts about which electrode does what rather than long calculations.

Concept 1 of 3: Galvanic and electrolytic cells compared

In both kinds of cell the two half-reactions happen at separate electrodes, and the electrons travel between them through a wire. In a galvanic cell the reaction runs by itself and the moving electrons are the electric current we use. In an electrolytic cell we pump electrons in with a supply to force a reaction that would not otherwise happen.

Definition

The anode is always where oxidation happens; the cathode is always where reduction happens. Memory aid: AN OX, RED CAT.
In a galvanic (voltaic) cell the anode is the negative electrode, because the oxidation there releases electrons into the wire.
In an electrolytic cell the anode is the positive electrode, because the supply pulls electrons out of it.
A salt bridge (or porous barrier) in a galvanic cell lets ions move between the two solutions so that the circuit is complete and each solution stays electrically neutral.
Electrons flow through the external wire; ions carry the current through the solution.

Feature	Galvanic (voltaic) cell	Electrolytic cell
Energy change	Chemical to electrical; the reaction is spontaneous	Electrical to chemical; the reaction is forced
Anode	Oxidation; negative electrode	Oxidation; positive electrode
Cathode	Reduction; positive electrode	Reduction; negative electrode
Electrons in the wire	From anode to cathode	Pushed by the supply into the cathode and drawn out of the anode
Examples	Daniell cell (zinc and copper), batteries, fuel cells	Electrolysis of brine, aluminium extraction, electroplating

Practice this conceptself-check · 3 quick reps

The anode's sign changes, its job does not

Oxidation always happens at the anode, and reduction at the cathode. But the anode is negative in a galvanic cell and positive in an electrolytic cell. Learn the electrodes by their reactions, not by their charges.

Concept 2 of 3: Standard reduction potentials and the cell voltage

Each half-cell has a tendency to take electrons, measured as its standard reduction potential, E°. Join two half-cells and the one with the higher E° wins: it is reduced, and it forces the other to be oxidised. The voltage of the cell is how far apart the two values are.

Definition

The standard reduction potential, E°, is measured against the standard hydrogen electrode, which is set at 0.00 V, at 25 °C with 1 mol/L solutions and 1 atm of gas.
The more positive E°, the stronger the oxidising agent on the left of the half-equation (it is reduced more easily). The more negative E°, the stronger the reducing agent on the right.
In a cell, the half-cell with the higher E° is the cathode; the lower one is the anode.
A positive cell potential means the reaction is spontaneous as written.
Useful values: $\mathrm{Mg^{2+}/Mg}$ $-2.37$ V, $\mathrm{Zn^{2+}/Zn}$ $-0.76$ V, $\mathrm{Fe^{2+}/Fe}$ $-0.44$ V, $\mathrm{H^+/H_2}$ 0.00 V, $\mathrm{Cu^{2+}/Cu}$ $+0.34$ V, $\mathrm{Ag^+/Ag}$ $+0.80$ V, $\mathrm{Cl_2/Cl^-}$ $+1.36$ V. The order matches the reactivity series.

Standard cell potential

E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

$E^\circ_{\text{cathode}}$ reduction potential of the half-cell that is reduced (the higher value)
$E^\circ_{\text{anode}}$ reduction potential of the half-cell that is oxidised

Worked example

A Daniell cell joins a zinc half-cell (

E^\circ = -0.76

V) to a copper half-cell (

E^\circ = +0.34

V). Find the anode, the cathode and the cell voltage.

Practice this conceptself-check · 4 quick reps

Never multiply E° by the equation's coefficients

E° is a property of the half-reaction, like a temperature, not an amount. Doubling a half-equation to balance electrons leaves its E° unchanged. Options that double or triple a potential are built from this mistake.

Concept 3 of 3: Electrolysis: what forms at each electrode

In electrolysis the supply drags electrons out of the anode and pushes them into the cathode. Positive ions travel to the cathode and are reduced; negative ions travel to the anode and are oxidised. In water there is competition: water itself can be reduced to hydrogen or oxidised to oxygen, and whichever needs less energy happens.

Definition

Molten ionic compound: the metal forms at the cathode, the non-metal at the anode.
Aqueous solution, cathode: a metal below hydrogen (Cu, Ag) is deposited; for a reactive metal (K, Na, Ca, Mg, Al) hydrogen forms instead: $\mathrm{2H_2O + 2e^- \rightarrow H_2 + 2OH^-}$ .
Aqueous solution, inert anode: a halide in concentrated solution gives the halogen: $\mathrm{2Cl^- \rightarrow Cl_2 + 2e^-}$ . Otherwise (sulfate, nitrate, dilute solutions) oxygen forms: $\mathrm{2H_2O \rightarrow O_2 + 4H^+ + 4e^-}$ .
Faraday's law in outline: the amount of product is proportional to the charge passed. One mole of electrons carries about 96 500 C (the Faraday constant), and charge = current × time.

Electrolyte (inert electrodes unless stated)	At the cathode (negative)	At the anode (positive)
Molten sodium chloride	Sodium metal	Chlorine
Concentrated aqueous sodium chloride (brine)	Hydrogen; sodium hydroxide is left in solution	Chlorine
Dilute sulfuric acid	Hydrogen	Oxygen, half the volume of the hydrogen
Aqueous copper(II) sulfate	Copper	Oxygen; the solution becomes acidic
Aqueous copper(II) sulfate, copper electrodes	Copper is deposited	The copper anode dissolves as $\mathrm{Cu^{2+}}$ (used to purify copper)
Molten aluminium oxide in cryolite	Aluminium	Oxygen, which burns the carbon anode to $\mathrm{CO_2}$

Practice this conceptself-check · 4 quick reps

A reactive metal is not deposited from its aqueous solution

Electrolysing aqueous sodium, potassium or magnesium salts gives hydrogen at the cathode, because water is reduced more easily than those ions. The metal itself is obtained only by electrolysing the molten compound.

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

Formulas (1)

Standard reduction potentials and the cell voltage
Standard cell potential
$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$

Reference tables (2)

Galvanic and electrolytic cells compared5 rows

Feature	Galvanic (voltaic) cell	Electrolytic cell
Energy change	Chemical to electrical; the reaction is spontaneous	Electrical to chemical; the reaction is forced
Anode	Oxidation; negative electrode	Oxidation; positive electrode
Cathode	Reduction; positive electrode	Reduction; negative electrode
Electrons in the wire	From anode to cathode	Pushed by the supply into the cathode and drawn out of the anode
Examples	Daniell cell (zinc and copper), batteries, fuel cells	Electrolysis of brine, aluminium extraction, electroplating

Electrolysis: what forms at each electrode6 rows

Electrolyte (inert electrodes unless stated)	At the cathode (negative)	At the anode (positive)
Molten sodium chloride	Sodium metal	Chlorine
Concentrated aqueous sodium chloride (brine)	Hydrogen; sodium hydroxide is left in solution	Chlorine
Dilute sulfuric acid	Hydrogen	Oxygen, half the volume of the hydrogen
Aqueous copper(II) sulfate	Copper	Oxygen; the solution becomes acidic
Aqueous copper(II) sulfate, copper electrodes	Copper is deposited	The copper anode dissolves as $\mathrm{Cu^{2+}}$ (used to purify copper)
Molten aluminium oxide in cryolite	Aluminium	Oxygen, which burns the carbon anode to $\mathrm{CO_2}$

Watch out for (3)

The anode's sign changes, its job does not
→ Galvanic and electrolytic cells compared
Never multiply E° by the equation's coefficients
→ Standard reduction potentials and the cell voltage
A reactive metal is not deposited from its aqueous solution
→ Electrolysis: what forms at each electrode

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