Electrolytic cells

1. Definition

• Electrolytic cells use electrical energy to drive a non-spontaneous redox reaction.
• They are commonly used in electroplating, refining metals, and the electrolysis of water.

2. How Electrolysis Works

• Electrolysis is the process where an electric current splits an ionic compound into its elements.
• Ions move towards electrodes:
  • Negative ions (anions) migrate to the anode, where they lose electrons (oxidation).
  • Positive ions (cations) migrate to the cathode, where they gain electrons (reduction).
• Electrons flow from the external circuit (battery) to the cathode.

3. Electrolysis of Molten Ionic Compounds

• In molten compounds (e.g., molten $NaCl$):

  • Anode reaction (oxidation): $2Cl^- \rightarrow Cl_2 + 2e^-$

  • Cathode reaction (reduction): $Na^+ + e^- \rightarrow Na$

4. Electrolysis of Ionic Solutions

• The electrolyte contains both the dissolved ionic compound and water.
• Rules for product formation:
  • At the anode (oxidation):
  • If the ionic compound contains a halide ion $(Cl⁻, Br⁻, I⁻),$ it will be oxidised to produce a halogen.
  • Otherwise, water is oxidised to produce oxygen gas.
  • At the cathode (reduction):
  • If the solution contains a less reactive metal (e.g., $Cu²⁺, Ag⁺, Au³⁺$), the metal will deposit at the cathode.
  • If the metal is more reactive than hydrogen (e.g., $Na⁺, K⁺, Ca²⁺$), H₂ gas is produced instead.

Example: Electrolysis of Copper(II) Chloride $(CuCl_2)$

• At the anode: (Chlorine gas forms.) $2Cl^- \rightarrow Cl_2 + 2e^-$

• At the cathode: (Copper metal deposits at the cathode.) $Cu^{2+} + 2e^- \rightarrow Cu$

• Overall reaction: $CuCl_2 \rightarrow Cu + Cl_2$

5. Electrolysis of Water

• Pure water is a poor conductor, so sulphuric acid $(H₂SO₄)$ is added to increase conductivity.

• Electrode Reactions:

  • At the anode (oxidation): $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$

  • At the cathode (reduction): $2H^+ + 2e^- \rightarrow H_2$

• Overall Reaction: $2H_2O \rightarrow 2H_2 + O_2$

6. Key Takeaways

• Electrolytic cells require an external power source to drive non-spontaneous reactions.
• Anode is positive (oxidation occurs), and cathode is negative (reduction occurs).
• Electrons flow from anode to cathode through the external circuit.
• Electrolysis is used in electroplating, refining metals, and hydrogen production.