Redox ReactionsMind Map

Oxidation and reduction were first defined using oxygen and hydrogen. Oxidation meant addition of oxygen or removal of hydrogen, while reduction meant removal of oxygen or addition of hydrogen. These classical definitions explain many reactions but fail for reactions without oxygen or hydrogen. The modern concept defines oxidation as loss of electrons and reduction as gain of electrons. Since electrons lost by one species must be gained by another, oxidation and reduction are always simultaneous. The species that causes oxidation is the oxidizing agent and itself gets reduced. The species that causes reduction is the reducing agent and itself gets oxidized. This topic is essential for identifying redox roles quickly.

Oxidation number is the apparent charge assigned to an atom in a molecule or ion by assuming complete transfer of bonding electrons to the more electronegative atom. It is not always the same as valency; valency shows combining capacity, while oxidation number can be positive, negative, zero, fractional or even average. Oxidation number rules help calculate oxidation states in elements, compounds and polyatomic ions. In redox reactions, the element whose oxidation number increases is oxidized and the element whose oxidation number decreases is reduced. Disproportionation occurs when the same element in one oxidation state undergoes both oxidation and reduction. NEET frequently tests rules, exceptions and quick calculations.

Balancing redox reactions requires both mass conservation and charge conservation. The oxidation number method balances total increase and decrease in oxidation numbers, then balances remaining atoms. The ion-electron method splits the reaction into oxidation and reduction half-reactions, balances atoms and charge separately, equalizes electrons, and adds the half-reactions. In acidic medium, oxygen is balanced using H₂O and hydrogen using H⁺. In basic medium, OH⁻ and H₂O are used, often by first balancing as acidic and then neutralizing H⁺ with OH⁻. NEET questions commonly test acidic/basic balancing, identification of oxidized and reduced species, and stoichiometric coefficients.

Redox reactions can be separated into oxidation and reduction half-reactions at electrodes. An electrode is a conducting surface where electron transfer occurs. Oxidation takes place at the anode and reduction takes place at the cathode. In a galvanic cell, a spontaneous redox reaction produces electrical energy; electrons flow through the external circuit from anode to cathode. Oxidation potential measures tendency to lose electrons, while reduction potential measures tendency to gain electrons. A species with higher reduction potential is more easily reduced and acts as a stronger oxidizing agent. Redox processes explain batteries, corrosion, extraction of metals, electrolysis, respiration and many analytical reactions.