Electrode Potentials & Half Cells | A-level Chemistry | OCR, AQA, Edexcel

Introduction

This tutorial explores electrode potentials and half-cells, fundamental concepts in electrochemistry. Understanding these topics is crucial for A-level chemistry students, as they are integral to redox reactions and the functioning of electrochemical cells. This guide will break down the key concepts and processes involved, making it easier for you to grasp these essential ideas.

Step 1: Understand Redox Reactions

• Redox reactions involve the transfer of electrons between substances.
• They consist of two half-reactions:
  • Oxidation: Loss of electrons.
  • Reduction: Gain of electrons.
• The flow of electrons generates electrical energy, which can be harnessed in electrochemical cells.

Step 2: Split Up Redox Reactions

• To analyze redox reactions, separate them into half-equations:
  1. Oxidation half-equation: Represents the generation of electrons.
  2. Reduction half-equation: Represents the consumption of electrons.
• This separation allows for the construction of half-cells, where the electron flow can be utilized.

Step 3: Construct Metal-Solution Half Cells

• A half-cell typically consists of a metal in contact with a solution of its ions.
• Key points:
  • Equilibrium is established between the metal and its ions.
  • The forward reaction involves the gain of electrons (reduction).
• Example: A copper electrode in a copper(II) sulfate solution.

Step 4: Create Half Cells with Gases

• When a gas is involved in the redox reaction, it can be bubbled through the solution:
  • This provides a surface for the reaction.
  • Example: Chlorine gas being reduced in a liquid medium.
• Important features of the electrode:
  • Inertness: The platinum electrode does not react with the substances.
  • Conductivity: It allows for electron transfer.
  • Surface area: Coating with platinum black increases the surface area for reactions.

Step 5: Set Up Half Cells with Aqueous Ions

• When both oxidation states are in aqueous form:
  • Use an equimolar solution and a platinum electrode.
  • Establish a dynamic equilibrium between the different oxidation states.

Step 6: Identify Which Half-Cell is Which

• To determine which half-cell releases electrons and which accepts them:
  • Analyze the equilibrium positions of the half-cell equations.
  • A left-shifted equilibrium indicates a stronger tendency to release electrons (oxidation).
  • A right-shifted equilibrium indicates a stronger tendency to accept electrons (reduction).
• The standard electrode potential (Eθ) quantifies this tendency.

Step 7: Use the Standard Hydrogen Electrode

• The standard hydrogen electrode serves as a reference point:
  • It is assigned a potential of 0V.
  • Connect it to other half-cells to measure their standard electrode potentials.
• Setup includes:
  • A salt bridge (often paper soaked in KNO3 or NH4NO3) to complete the circuit.
  • A high-resistance voltmeter to measure voltage differences.

Conclusion

Understanding electrode potentials and half-cells is vital for studying electrochemistry. By following these steps, you can effectively analyze redox reactions and their applications in electrochemical cells. For further study, consider practicing problems related to electrode potentials and exploring real-world applications of electrochemistry in batteries and corrosion.