A modular guide to driving non-spontaneous redox reactions with electrical energy

Electrolysis lies at the heart of industrial chemistry, materials processing, and sustainable energy systems. Whereas galvanic cells convert chemical energy into electrical energy, electrolytic cells do the reverse: they use an external electric current to force a non-spontaneous chemical reaction.

This lecture explores the theory and operation of electrolytic cells, examines electrode processes and energy requirements, and introduces practical examples such as water electrolysis, electroplating, and metal extraction. We also discuss overpotential, activation energy, and cell efficiency, key factors determining how effectively electrical energy is converted into chemical change.

1. What Is Electrolysis?

Electrolysis is a chemical process driven by electricity. When an electric current passes through an electrolyte, an ionic substance dissolved, or molten ions, move towards electrodes where oxidation and reduction occur.

A simple definition:

Electrolysis is the decomposition of a compound by the passage of an electric current.

The overall principle obeys Faraday’s laws of electrolysis, first established by Michael Faraday (1834), which state that the mass of substance liberated at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

Mathematically:

where:

m = mass of substance liberated (g)

Q = total charge passed (C)

M = molar mass (g/mol)

n = number of electrons transferred per ion

F = Faraday constant = 96,485 C/mol

Reference: Royal Society of Chemistry – Faraday’s Laws of Electrolysis

2. Components of an Electrolytic Cell

An electrolytic cell comprises several essential parts:

• Power source – provides the external current (usually DC)
• Electrodes – conductors where oxidation and reduction occur
• Electrolyte – ionic medium enabling charge transfer
• Ions – carry electric current through the electrolyte

Electrode polarity

• The anode is positive in electrolytic cells (connected to the positive terminal of the power supply). Oxidation occurs here.
• The cathode is negative (connected to the negative terminal). Reduction occurs here.

This is opposite to galvanic cells, where oxidation occurs at the negative electrode.

3. How Electrolysis Works: The Redox Perspective

During electrolysis, the external power source pushes electrons into the cathode and pulls electrons from the anode, forcing ions in the electrolyte to migrate:

• Cations (+) move towards the cathode to gain electrons (reduction).
• Anions (-) move towards the anode to lose electrons (oxidation).

Example: Electrolysis of molten sodium chloride

Half-reactions:

• Cathode (reduction):

• Anode (oxidation):

This process produces metallic sodium and chlorine gas, an essential industrial method known as the Downs process.

Reference: Chemguide – Electrolysis of Molten Sodium Chloride

4. Electrolysis of Aqueous Solutions

In aqueous systems, both water and solute ions may be reduced or oxidised. The products depend on electrode potentials and concentrations.

Example: Electrolysis of Aqueous Sodium Chloride (Brine)

Overall reaction:

Half-reactions:

Cathode (reduction):

Anode (oxidation):

The overall reaction yields hydrogen gas, chlorine gas, and sodium hydroxide, forming the basis of the chlor-alkali industry.

Reference: LibreTexts – Electrolysis of Aqueous Solutions

5. Overvoltage and Activation Energy

In practice, the applied potential required to drive an electrolytic reaction is often greater than the theoretical decomposition voltage calculated from standard potentials. This difference is known as overpotential or overvoltage, arising from:

• Kinetic barriers (activation energy for electron transfer)
• Surface effects (electrode material, roughness, adsorbed species)
• Mass-transport limitations (diffusion of reactants/products)

Example: Water electrolysis

Overall reaction:

The theoretical decomposition potential of water is 1.23 V, but in practice, 1.6–2.0 V is needed.

This higher voltage is required because of overpotentials at both electrodes, additional energy losses due to kinetic barriers, gas bubble formation, and electrode surface effects.

Half-reactions:

Anode (oxidation):

Cathode (reduction):

Explanation:

During electrolysis, water is decomposed into hydrogen and oxygen gases:

• At the anode, water molecules are oxidised to oxygen gas, releasing protons and electrons.
• At the cathode, these protons gain electrons to form hydrogen gas.

This process is non-spontaneous, meaning it requires an external electrical energy source to proceed.

The overpotential (extra voltage) arises from factors such as:

• Electrode kinetics (slow reaction rates)
• Gas evolution barriers (bubbles hinder active surface area)
• Electrode material effects (some surfaces facilitate charge transfer better than others)

Efficient electrocatalysts like platinum, iridium oxide, or nickel-based materials are used to lower these overpotentials and improve the efficiency of hydrogen production.

Reference: Nature Energy – Water Splitting Catalysis

6. Energy and Efficiency

The electrical work done (W elec) to drive electrolysis is:

where
= electrons per mole of reaction
= Faraday constant
= applied cell voltage

The efficiency of an electrolytic cell depends on the ratio of theoretical to actual energy consumption:

Losses arise from resistance (IR drop), overpotential, and heat generation.

7. Industrial Applications of Electrolysis

Electrolysis has vast technological importance across multiple sectors:

a) Electrorefining and Electrowinning

Used to purify metals or extract them from ores.

• Copper refining:

Impure copper dissolves, and pure copper plates out, leaving behind impurities.

Reference: Metallurgy of Copper – Refining Process (RSC)

b) Electroplating

Deposition of a thin metal layer to protect or beautify a surface.

Example: Silver plating a spoon.

Electroplating improves corrosion resistance, aesthetic appeal, and conductivity.

c) Production of Chemicals

• Chlor-alkali process: Production of chlorine (Cl₂), hydrogen (H₂), and sodium hydroxide (NaOH).
• Hydrogen generation: key step in ammonia synthesis (Haber–Bosch process).
• Electrosynthesis: formation of organic compounds using electricity instead of reagents.

8. Factors Affecting Electrolysis

Several variables influence product distribution and reaction rate:

Factor	Effect
Nature of electrodes	Inert electrodes (Pt, C) vs. reactive (Cu, Fe) determine reaction pathways
Electrolyte concentration	Competing reactions depend on ion availability
Applied potential/current	Determines which redox reactions occur
Temperature	Increases ion mobility and reaction kinetics
pH	Alters redox equilibria, especially in aqueous systems

Example: Electrolysis of dilute H₂SO₄ primarily produces hydrogen and oxygen gases through the decomposition of water. In concentrated acid, side reactions become significant.

9. Faraday’s Laws in Practice

First Law

The mass (m) of a substance produced at an electrode is proportional to the total charge passed.

where Z is the electrochemical equivalent.

Second Law

The masses of different substances liberated by the same charge are proportional to their chemical equivalent weights.

Passing 96,485 C (1 Faraday) through molten AgNO₃ deposits 107.9 g of Ag, calculated as molar mass 107.9 g/mol ÷ 1 electron per ion = 107.9 g.

Passing the same charge through molten CuSO₄ deposits 31.75 g of Cu, calculated as molar mass 63.5 g/mol ÷ 2 electrons per ion = 31.75 g.

10. Overpotential and Kinetic Limitations

Overpotential stems from electrode kinetics. To visualise this, the Butler–Volmer equation expresses the current density (i) as a function of overpotential (η):

where is the exchange current density, a measure of electrode activity.

High i₀ and low η indicate a fast reaction (efficient electrode).

Catalytic surfaces such as Pt, IrO₂, or NiFe alloys are engineered to reduce η and improve performance.

Reference: Electrochemical Kinetics – Butler–Volmer Model (LibreTexts)

11. Electrolysis and the Energy Transition

In the 21st century, electrolysis plays a pivotal role in green energy:

• Hydrogen production from renewable electricity enables carbon-neutral fuel cycles.
• Power-to-gas technologies store intermittent solar and wind energy.
• CO₂ electroreduction converts carbon dioxide into synthetic fuels and feedstocks.
• Electrochemical recycling recovers valuable metals from waste streams.

Advances in membrane electrolyser design (PEM, solid-oxide, alkaline) are pushing efficiencies toward 80–90 %.

Reference: International Energy Agency – The Future of Hydrogen (2019)

12. Laboratory Demonstration: Copper(II) Sulphate Electrolysis

In a school or university laboratory, electrolysis of aqueous copper(II) sulphate illustrates key principles:

Half-Reactions

Cathode (Reduction):
At the cathode, copper(II) ions are reduced to metallic copper:

Anode (Oxidation):
At the anode, water is oxidised to produce oxygen gas and hydrogen ions:

2 H₂O(l) → O₂(g) + 4 H⁺(aq) + 4 e⁻

If copper electrodes are used, both anode dissolution and cathode deposition occur, maintaining the solution concentration. This serves as a model for industrial copper refining.

Observations:

• The blue color fades if an inert anode is used (Cu²⁺ is consumed).
• Gas bubbles (O₂) form at the anode.

13. Quantitative Example

Problem:
How much aluminum (in grams) can be produced by passing 96,485 C through molten Al₂O₃?

Reaction at the cathode:
Al³⁺ + 3 e⁻ → Al(s)

Given:

Faraday constant, F = 96,485 C/mol

n = 3 (electrons per Al³⁺)

Molar mass of Al, M = 26.98 g/mol

So, 1 F deposits 8.99 g of aluminium.

14. Limitations and Environmental Considerations

While electrolysis enables cleaner processes, it also demands significant electrical input. Sustainability depends on renewable power sources and efficient electrolyser technologies.

Challenges include:

• High capital cost of catalysts and membranes.
• Energy losses via heat and overpotential.
• Resource demands for electrode materials (e.g. iridium, platinum).

Emerging solutions focus on earth-abundant catalysts (e.g. Ni, Co, Fe oxides) and modular flow systems for large-scale hydrogen and ammonia synthesis.

Reference: Royal Society – Green Ammonia and Hydrogen

15. Summary

Electrolysis exemplifies how electricity can be used to transform matter. Key takeaways:

1. Electrolytic cells convert electrical energy into chemical energy, driving non-spontaneous redox reactions.
2. Oxidation occurs at the anode (positive); reduction occurs at the cathode (negative).
3. Faraday’s laws relate the charge passed to the amount of substance formed.
4. Overvoltage and activation energy influence efficiency.
5. Industrial applications include metal extraction, electroplating, chemical synthesis, and hydrogen production.
6. Modern electrolysis underpins green energy technologies vital for a sustainable future.

Further Reading and References

• Chemguide – Electrolysis
• LibreTexts – Electrolytic Cells
• Nature Energy – Advances in Electrolysis
• Royal Society of Chemistry – Faraday’s Laws
• International Energy Agency – Hydrogen Report

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