An electrolytic cell is a type of electrochemical cell that uses electrical energy to force a non-spontaneous chemical reaction to occur. It is the opposite of a galvanic (voltaic) cell. Electrolytic cells are essential in IB Chemistry Topic 9 (Redox Processes) and Topic 19 (HL), particularly in the contexts of electroplating, metal extraction, and electrolysis of molten or aqueous ionic compounds. Understanding electrolytic cells strengthens your knowledge of electron flow, electrode roles, and redox behavior.
What Is an Electrolytic Cell?
An electrolytic cell is a system that uses an external power source to drive a chemical reaction that would not occur on its own.
Key features:
- Electrical energy → chemical change
- Requires a power supply (battery/DC source)
- Redox reactions are forced
- Electrons are pushed in the opposite direction compared to a galvanic cell
Electrolytic cells are widely used in industry to produce chemicals, extract metals, and refine materials.
How an Electrolytic Cell Works
An electrolytic cell contains:
- Two electrodes (anode and cathode)
- An ionic substance (molten or aqueous)
- A power source to force electron movement
1. Power supply forces electron flow
Electrons are pushed from the positive terminal toward the negative terminal.
2. Oxidation occurs at the anode
Anode is positive in an electrolytic cell because it attracts anions.
3. Reduction occurs at the cathode
Cathode is negative because it supplies electrons.
Even though electrode signs change, oxidation always occurs at the anode and reduction always at the cathode.
Electron Flow and Ion Movement
Electron flow:
Power source → cathode → anode
Ion movement:
- Cations move to the cathode (reduction)
- Anions move to the anode (oxidation)
This movement completes the circuit and allows redox reactions to occur.
Example 1: Electrolysis of Molten NaCl
Molten NaCl contains only Na⁺ and Cl⁻ ions.
Cathode (reduction):
Na⁺ + e⁻ → Na(l)
Anode (oxidation):
2Cl⁻ → Cl₂(g) + 2e⁻
Products:
- Liquid sodium metal
- Chlorine gas
This process is widely used in metal extraction.
Example 2: Electrolysis of Aqueous Solutions
Aqueous solutions include water, so multiple species may compete for oxidation or reduction.
Example: Aqueous CuSO₄ with inert electrodes
Cathode:
Cu²⁺ + 2e⁻ → Cu(s)
Copper plates onto the cathode.
Anode:
2H₂O → O₂(g) + 4H⁺ + 4e⁻
Water is oxidized instead of sulfate.
This demonstrates why electrolysis of solutions requires understanding of electrode potentials.
Example 3: Electroplating
Electrolytic cells can coat a metal object with another metal.
Procedure:
- Object to be plated = cathode
- Metal source = anode
- Solution contains ions of the plating metal
Example:
Electroplating silver onto a spoon using Ag⁺ solution.
This technique is common in jewelry, electronics, and corrosion protection.
Why Electrolytic Cells Matter in IB Chemistry
1. Metal extraction
Electrolysis extracts reactive metals like aluminum and sodium.
2. Industrial chemical production
Such as:
- Chlorine gas
- Hydrogen gas
- Sodium hydroxide
3. Refining metals
Copper is purified through electrolytic refining.
4. Electroplating
Used to coat objects with protective or decorative metals.
5. Redox understanding
Electrolytic cells help students visualize oxidation and reduction under forced conditions.
Galvanic vs Electrolytic Cells
Feature Galvanic Cell Electrolytic Cell Reaction Spontaneous Non-spontaneous Energy Produces electricity Consumes electricity Anode Negative Positive Cathode Positive Negative Electron flow Anode → cathode Power source → cathode
Understanding the differences is crucial for exam success.
Common IB Misunderstandings
“Cathode is always positive.”
False—cathode is positive in galvanic cells, negative in electrolytic cells.
“Electrolysis occurs on its own.”
No—external power is required.
“Water never reacts during electrolysis.”
Water often competes with ions in aqueous solutions.
“Electrodes must be metals.”
Graphite and platinum are common inert electrodes.
FAQs
Why is the anode positive in electrolysis?
Because the power supply pulls electrons away, forcing oxidation.
Does electrolysis always produce gases?
Not always—it can produce metals or ions depending on the system.
Can electrolysis occur at room temperature?
Yes, if the solution or molten salt allows ion movement.
Conclusion
An electrolytic cell uses electrical energy to drive non-spontaneous redox reactions. With a positive anode, a negative cathode, and forced electron flow, electrolytic cells are essential to metal extraction, electroplating, and industrial chemical production. Understanding electrolytic cells is critical for mastering redox chemistry in the IB syllabus.
