🔄 Redox Reactions

1. Oxidation and Reduction

Classical concept: Oxidation is addition of oxygen/electronegative element or removal of hydrogen/electropositive element. Reduction is the opposite.

⚡ Electronic Concept

Oxidation: Loss of electrons. (LEO / OIL)

Reduction: Gain of electrons. (GER / RIG)

Redox Reactions: Reactions where oxidation and reduction occur simultaneously.

Oxidising Agent (Oxidant): Accepts electrons (gets reduced).
Reducing Agent (Reductant): Donates electrons (gets oxidized).

2. Oxidation Number

The residual charge which an atom appears to have when all other atoms are removed from it as ions. It helps in keeping track of electron shifts.

Rules for Oxidation Number

1. Elements in free state = 0 (e.g., O₂, Na).

2. Fluorine is always -1 in its compounds.

3. Oxygen is usually -2 (except in peroxides where it is -1, and OF₂ where it is +2).

4. Hydrogen is +1 with non-metals and -1 with metals (hydrides).

5. The algebraic sum of oxidation numbers of all atoms in a neutral molecule is zero.

3. Types of Redox Reactions

Combination Reactions
Decomposition Reactions
Displacement Reactions (Metal and Non-metal)
Disproportionation Reactions: A reaction in which an element in one oxidation state is simultaneously oxidized and reduced. (e.g., 2H₂O₂ → 2H₂O + O₂).

4. Balancing Redox Reactions

Two main methods are used to balance chemical equations for redox processes:

Oxidation Number Method: Based on the fact that total increase in oxidation number equals total decrease in oxidation number.
Half-Reaction Method (Ion-Electron Method): The reaction is broken into two half-reactions (oxidation half and reduction half), balanced separately, and then added.

5. Redox Reactions and Electrode Processes

The basis of Electrochemical Cells (Galvanic Cells). The electrode where oxidation occurs is the Anode (negative). The electrode where reduction occurs is the Cathode (positive).