Redox Reactions and Electrochemistry – JEE Mains Chemistry

• by Zenith LMS
• Updated April 2025

1. Electronic Concepts of Oxidation and Reduction

Oxidation involves the loss of electrons, while reduction involves the gain of electrons. In a redox reaction, one species is oxidized, and another is reduced. These reactions are key to understanding various chemical processes.

2. Redox Reactions

A redox reaction is a chemical reaction in which the oxidation states of atoms change. The substance that gains electrons is reduced, and the substance that loses electrons is oxidized. These reactions are common in electrochemical cells.

3. Oxidation Number

• Oxidation Number: The charge an atom would have if it gained or lost electrons in a chemical reaction.
• Rules for Assigning Oxidation Number: The oxidation number of an element in its elemental form is zero. Oxygen usually has an oxidation number of -2, hydrogen is +1, etc.
• Balancing Redox Reactions: Balancing requires ensuring that the number of electrons lost equals the number of electrons gained in a reaction.

4. Electrolytic and Metallic Conduction

• Electrolytic Conduction: The conduction of electricity through an electrolyte solution by the movement of ions.
• Metallic Conduction: The conduction of electricity through metals, where electrons move freely between atoms.

5. Conductance in Electrolytic Solutions

• Molar Conductivity: The conductance of an electrolyte solution per mole of solute.
• Variation with Concentration: Molar conductivity decreases as concentration increases due to ion pairing and other factors.
• Kohlrausch’s Law: At infinite dilution, the molar conductivity of an electrolyte is constant and independent of concentration. This is used to determine the conductivity of weak electrolytes.

6. Electrochemical Cells

• Electrolytic Cells: Cells in which non-spontaneous reactions occur, requiring an external power source.
• Galvanic Cells: Cells that generate electricity from spontaneous reactions, such as a battery.
• Electrode Potential: The potential difference between an electrode and its ion in solution. It is a measure of the tendency of an electrode to gain or lose electrons.
• Standard Electrode Potential: The electrode potential measured under standard conditions (1 M concentration, 1 atm pressure, 25°C).
• Half-Cell Reactions: Each half of the electrochemical reaction, where oxidation occurs at one electrode, and reduction occurs at the other.
• Cell Reactions: The overall reaction in an electrochemical cell, combining the oxidation half-cell and the reduction half-cell.

7. Measurement of emf of Galvanic Cell

The emf (electromotive force) of a Galvanic cell is measured by connecting the two half-cells through a voltmeter. It indicates the cell's ability to produce an electric current.

8. Nernst Equation

The Nernst equation relates the cell potential to the concentrations of the ions involved in the redox reaction:E = E° - (RT/nF) ln Qwhere E is the cell potential, E° is the standard cell potential, R is the gas constant, T is temperature, n is the number of moles of electrons involved, F is the Faraday constant, and Q is the reaction quotient.

9. Relationship Between Cell Potential and Gibbs' Energy Change

The relationship is given by the equation:ΔG = -nFEwhere ΔG is the Gibbs free energy change, n is the number of electrons, F is Faraday’s constant, and E is the cell potential.

10. Types of Electrochemical Cells

• Dry Cell: A simple voltaic cell with a paste of electrolyte, commonly used in batteries.
• Lead Accumulator: A type of rechargeable battery used in vehicles.
• Fuel Cells: Electrochemical cells that convert the chemical energy of fuel (often hydrogen) directly into electricity.