Electrolysis
Objectives:
- Distinguish between electrolytes and non‑electrolytes.
- State and apply Faraday’s laws of electrolysis.
- Solve numerical problems using the formula: m = (Q × M) / (F × z).
- Explain the electrolysis reactions of dilute H₂SO₄, aqueous CuSO₄, aqueous CuCl₂, aqueous NaCl, and fused NaCl.
- Identify factors affecting the discharge of ions at the electrodes.
- Specify the applications of electrolysis such as purification of metals and production of elements/compounds (e.g. Al, Na, O₂, Cl₂, NaOH).
- Identify various electrochemical cells including the electrochemical series, half‑cell reactions, and electrode potentials.
- Describe corrosion as an electrolytic process and methods for protecting metals (cathodic protection, painting, electroplating, coating with grease/oil).
Key Concepts and Formulas
Electrolyte: A substance that conducts electricity when dissolved or molten because it produces ions.
Example: NaCl dissolves in water to produce Na⁺ and Cl⁻ ions.
Non-Electrolyte: A substance that does not conduct electricity in solution because it does not produce ions.
Example: Sugar in water does not produce ions and does not conduct electricity.
Faraday’s First Law: The mass (m) of a substance deposited is directly proportional to the total electric charge (Q) passed.
Formula: m = (Q × M) / (F × z)
Example: If Q = 2.0×10⁵ C, M (Cu) = 63.55 g/mol, z = 2, and F ≈ 96485 C/mol, then m ≈ 0.66 g of copper is deposited.
Faraday’s Second Law: For the same quantity of electricity, the masses of different substances deposited are proportional to their equivalent weights.
Example: Under the same charge, a metal with a lower equivalent weight deposits a larger mass.
Factors Affecting Ion Discharge: The discharge of ions depends on ion concentration, electrode material, overpotential, and temperature.
Example: Increasing NaCl concentration enhances conductivity, leading to higher Cl₂ evolution at the anode.
Electrolysis Reactions – Part I
This page explains typical electrolysis reactions with worked examples immediately following each description.
Electrolysis of Dilute H₂SO₄:
Reaction: 2H₂SO₄ (aq) → 2H₂ + O₂ + 2SO₄²⁻
Explanation: In dilute H₂SO₄, water is decomposed. H⁺ ions gain electrons at the cathode to form H₂ gas; water is oxidized at the anode to form O₂ gas.
Worked Example: Observe gas evolution at both electrodes during the electrolysis of dilute H₂SO₄.
Electrolysis of Aqueous CuSO₄:
Reaction: CuSO₄ (aq) → Cu (deposited at cathode) + O₂ (anode) + SO₄²⁻ remains
Explanation: Cu²⁺ ions gain electrons at the cathode to form copper metal, while water oxidizes at the anode to produce oxygen.
Worked Example: In a CuSO₄ cell, copper plates out on the cathode while oxygen is seen at the anode.
Electrolysis of Aqueous CuCl₂:
Reaction: CuCl₂ (aq) → Cu (cathode) + Cl₂ (anode)
Explanation: Copper ions are reduced to form copper at the cathode, and chloride ions are oxidized to form chlorine gas at the anode.
Worked Example: The production of chlorine gas, with its characteristic odor, confirms the reaction.
Electrolysis Reactions – Part II & Numerical Applications
This page covers additional reactions and numerical applications.
Electrolysis of Aqueous NaCl:
Reaction: 2NaCl (aq) + 2H₂O → 2NaOH + Cl₂ + H₂
Explanation: Both water and chloride ions participate. At the cathode, water is reduced to form NaOH and H₂, while Cl⁻ ions are oxidized at the anode to produce Cl₂ gas.
Worked Example: The alkaline nature of the solution confirms NaOH production, and chlorine gas is detected at the anode.
Electrolysis of Fused NaCl:
Reaction: 2NaCl (molten) → 2Na + Cl₂
Explanation: In molten NaCl, water is absent. Sodium ions are reduced at the cathode to form sodium metal, while chloride ions are oxidized at the anode to form chlorine gas.
Worked Example: This process is used industrially to produce pure sodium metal and chlorine gas.
Example 6 – Numerical Calculation Using Faraday’s First Law:
Given: Q = 2.0×10⁵ C, M (Cu) = 63.55 g/mol, z = 2, F ≈ 96485 C/mol.
Calculation: m = (2.0×10⁵ C × 63.55 g/mol) / (96485 C/mol × 2) ≈ 0.66 g
Explanation: This calculation shows how to apply Faraday’s law to determine the mass of copper deposited during electrolysis.
Example 7 – Effect of Ion Concentration:
Explanation: Increasing the concentration of NaCl increases the number of ions available, thereby increasing the conductivity and the deposition rate of chlorine at the anode under constant current.
Worked Example: A 1 M NaCl solution deposits more mass than a 0.1 M solution under identical conditions.
Example 8 – Effect of Electrode Material:
Explanation: The electrode material affects the overpotential. Platinum electrodes in aqueous CuSO₄ typically yield a higher copper deposition compared to graphite electrodes.
Worked Example: Experiments show that platinum electrodes produce a higher mass of copper deposition due to lower overpotential.
Applications of Electrolysis, Electrochemical Cells & Corrosion Protection
This page covers the practical applications of electrolysis, the fundamentals of electrochemical cells, and methods of corrosion protection.
Uses of Electrolysis:
Electrolysis is employed for the purification of metals (e.g., copper), production of elements and compounds (e.g., aluminum, sodium, oxygen, chlorine, sodium hydroxide), and in various industrial processes.
Worked Example: The electrolytic refining of copper involves using impure copper as an anode and pure copper as a cathode, resulting in high-purity copper deposition.
Electrochemical Cells:
An electrochemical cell consists of two half-cells. The electrochemical series (e.g., K, Ca, Na, Mg, Al, Zn, Fe, Sn, Pb, H, Cu, Hg, Ag, Au) ranks elements by their electrode potentials.
Key Concept: Half-cell reactions and the Nernst equation are used to calculate electrode potentials (simple calculations only).
Worked Example: In a Cu/Cu²⁺ half-cell, the standard electrode potential is +0.34 V.
Corrosion as an Electrolytic Process:
Corrosion involves the oxidation of metals (e.g., iron rusting). Cathodic protection, painting, electroplating, and coating with grease or oil are methods used to prevent corrosion.
Worked Example: Cathodic protection is used in pipelines where a sacrificial anode (e.g., zinc) is used to prevent the oxidation of the main metal.
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