Introduction:

Redox reactions (oxidation-reduction reactions) play a crucial role in chemistry, from industrial applications to biological systems. This topic is fundamental in electrochemistry, influencing processes like corrosion, energy storage, and electrode potentials. Understanding redox reactions and electrode processes helps students grasp essential concepts for competitive exams and practical applications.



Redox Reactions – An Overview:

Oxidation and Reduction

- Oxidation: Loss of electrons or increase in oxidation state.

- Reduction: Gain of electrons or decrease in oxidation state.

Oxidizing and Reducing Agents:

- Oxidizing agent: Gains electrons and gets reduced.

- Reducing agent: Loses electrons and gets oxidized.

Types of Redox Reactions:

- Combination Reactions: Two or more reactants form a single product.

- Decomposition Reactions: A compound breaks down into simpler substances.

- Displacement Reactions: One element replaces another in a compound.

- Disproportionation Reactions: The same element undergoes oxidation and reduction simultaneously.

Oxidation Number and Balancing Redox Reactions:

Oxidation Number Rules:

- The oxidation number of free elements is zero.

- Oxygen generally has an oxidation number of -2 (except in peroxides and superoxides).

- Hydrogen has an oxidation number of +1 (except in metal hydrides where it is -1).

Balancing Redox Reactions:

- Oxidation Number Method: Identify oxidation state changes and balance electrons.

- Ion-Electron Method: Divide reactions into oxidation and reduction half-reactions and balance them separately.

Read Also: Temperature Dependence of the Rate of a Reaction - Class 12 Chemistry Notes

Electrode Processes and Electrochemical Cells:

Types of Electrochemical Cells:

- Galvanic (Voltaic) Cells: Convert chemical energy into electrical energy (e.g., Daniel cell).

- Electrolytic Cells: Use electrical energy to drive non-spontaneous reactions (e.g., electrolysis of water).

Standard Electrode Potential (E°):

- The potential difference between a metal electrode and its ion solution under standard conditions.

- Measured in volts and compared to the standard hydrogen electrode (SHE).

Nernst Equation:

- The Nernst equation calculates the electrode potential of a cell under non-standard conditions:

E = E° - (RT/nF) ln Q

Where:

- E = electrode potential,

- E° = standard electrode potential,

- R = universal gas constant,

- T = temperature,

- n = number of electrons transferred,

- F = Faraday’s constant,

- Q = reaction quotient.

Applications of Redox Reactions and Electrode Processes:

Industrial Applications:

- Metallurgy: Extraction of metals using redox processes.

- Corrosion Prevention: Galvanization and cathodic protection.

- Electroplating: Coating metals to prevent rusting.

Biological Importance:

- Cellular Respiration: Energy production in living organisms.

- Photosynthesis: Redox reactions drive glucose synthesis in plants.

Batteries and Fuel Cells:

- Lithium-ion Batteries: Power electronic devices.

- Hydrogen Fuel Cells: Alternative energy sources.

Conclusion:

Redox reactions and electrode processes are integral to chemistry and real-world applications. Mastering oxidation-reduction concepts is crucial for understanding electrochemical cells, energy storage, and various industrial processes. A strong grasp of these principles will help students excel in their exams and practical applications in science and technology.