Electrochemistry: Galvanic and Electrolytic Cells

What Is Electrochemistry?

Electrochemistry is the study of the relationship between chemical reactions and electrical energy. It's the science behind batteries, corrosion, electroplating, and industrial production of metals like aluminium.

There are two types of electrochemical cells:

Galvanic (voltaic) cells: Spontaneous redox reactions generate electricity.
Electrolytic cells: External electricity drives non-spontaneous reactions.

Learning Goals: By the end of this guide, you should be able to:

Describe the components of galvanic and electrolytic cells.

Use standard electrode potentials to predict cell EMF.

Identify anode and cathode in both cell types.

Apply Faraday's laws to calculate quantities in electrolysis.

Galvanic (Voltaic) Cells

A galvanic cell converts chemical energy → electrical energy using a spontaneous redox reaction.

Key Components

Component	Function
Anode (–)	Oxidation occurs (electrons leave)
Cathode (+)	Reduction occurs (electrons arrive)
Salt bridge	Completes the circuit with ion flow; maintains electrical neutrality
External wire	Electrons flow from anode to cathode

Memory trick: Anode = Away = oxidation (An Ox). Cathode = Comes = reduction (Red Cat).

Cell EMF (Electromotive Force)

E°_{cell} = E°_{cathode} - E°_{anode}

A positive $E°_{cell}$ means the reaction is spontaneous.

Standard Electrode Potentials

Measured against the Standard Hydrogen Electrode (SHE, $E° = 0.00\ V$ ).

Half-Cell	$E°$ / V
$Li^+ + e^- \rightleftharpoons Li$	–3.04
$Zn^{2+} + 2e^- \rightleftharpoons Zn$	–0.76
$2H^+ + 2e^- \rightleftharpoons H_2$	0.00
$Cu^{2+} + 2e^- \rightleftharpoons Cu$	+0.34
$Ag^+ + e^- \rightleftharpoons Ag$	+0.80
$F_2 + 2e^- \rightleftharpoons 2F^-$	+2.87

More positive $E°$ → stronger oxidising agent (easier to reduce).
More negative $E°$ → stronger reducing agent (easier to oxidise).

Electrolytic Cells

An electrolytic cell uses electrical energy → chemical energy, forcing non-spontaneous reactions to occur.

Differences from Galvanic

Feature	Galvanic	Electrolytic
Spontaneity	Spontaneous	Non-spontaneous
Energy conversion	Chemical → Electrical	Electrical → Chemical
Anode charge	Negative (–)	Positive (+)
Cathode charge	Positive (+)	Negative (–)
Applications	Batteries, fuel cells	Electroplating, metal extraction

In electrolysis, the anode is connected to the positive terminal of the power supply, and the cathode to the negative terminal. Oxidation still occurs at the anode, reduction at the cathode.

Electrochemistry Cell Simulator

Build galvanic and electrolytic cells interactively. Choose electrode metals, see electron flow, and calculate cell EMF in real time.

Build a Cell

Faraday's Laws of Electrolysis

First Law

The mass of substance deposited at an electrode is directly proportional to the quantity of charge passed.

Second Law

The mass deposited is proportional to the molar mass and inversely proportional to the number of electrons transferred.

m = \frac{MIt}{nF}

Where: $M$ = molar mass, $I$ = current (A), $t$ = time (s), $n$ = electrons transferred, $F$ = Faraday constant ( $96485\ C\ mol^{-1}$ ).

Worked Examples

Example 1: Calculate Cell EMF

Question: What is the EMF of a Zn-Cu cell?

$E°_{cell} = E°_{cathode} - E°_{anode} = (+0.34) - (-0.76) = +1.10\ V$

Since $E°_{cell} > 0$ , the reaction is spontaneous. Zn is oxidised (anode), Cu²⁺ is reduced (cathode).

Example 2: Faraday's Law Calculation

Question: How much copper is deposited when 2.0 A flows for 30 minutes through $CuSO_4$ solution?

m = \frac{M \times I \times t}{n \times F} = \frac{63.5 \times 2.0 \times 1800}{2 \times 96485} = 1.18\ g

Example 3: Predicting Products of Electrolysis

Question: What products form at each electrode when molten $NaCl$ is electrolysed?

Cathode (–): $Na^+ + e^- \rightarrow Na$ (sodium metal deposited)
Anode (+): $2Cl^- \rightarrow Cl_2 + 2e^-$ (chlorine gas evolved)

Common Mistakes

Mixing up anode/cathode signs — In galvanic cells: anode is (–), cathode is (+). In electrolytic cells: it's reversed. But oxidation ALWAYS occurs at the anode regardless.
Subtracting $E°$ values incorrectly — Always: $E°_{cell} = E°_{cathode} - E°_{anode}$ . Don't reverse the sign of the more negative electrode arbitrarily.
Forgetting to convert time to seconds — In Faraday's law, time must be in seconds, not minutes.
Not considering aqueous solutions in electrolysis — When electrolyzing aqueous solutions (not molten), water can be oxidised or reduced preferentially depending on discharge potentials.

Exam Tips (A-Level / AP / IB)

Always label your diagram: anode (oxidation), cathode (reduction), direction of electron flow, ion movement in salt bridge.
For EMF calculations, identify which half-cell has the more positive (or less negative) $E°$ — that's the cathode.
In electrolysis questions, state what happens at EACH electrode with half-equations.

Frequently Asked Questions

What is the difference between a galvanic and electrolytic cell?

A galvanic cell uses a spontaneous redox reaction to produce electricity (like a battery). An electrolytic cell uses electricity to force a non-spontaneous reaction (like electroplating or extracting metals).

Why is the salt bridge necessary?

Without a salt bridge, charge would build up in each half-cell (too many positive ions in one, too many negative in the other), stopping the reaction. The salt bridge allows ions to flow, maintaining electrical neutrality.

What does a negative E°cell mean?

A negative $E°_{cell}$ means the reaction is non-spontaneous under standard conditions. You would need to supply external energy (electrolysis) to make it happen.

Balancing Redox Reactions — The half-equations used in electrochemistry.
Chemical Bonds — Ionic bonding in electrolytes.
Periodic Trends — Electrode potential trends relate to atomic structure.