Redox ReactionsNEET Chemistry · Class 11 · NCERT Chapter 7

Use these rules in order: (1) Free elements have oxidation state 0. (2) Monatomic ions have OS = their charge. (3) In compounds, F is always −1. (4) O is usually −2 (except in peroxides: −1, in OF₂: +2). (5) H is +1 with non-metals and −1 with metals. (6) The sum of all oxidation states equals the charge of the species (0 for neutral molecules, equal to ionic charge for ions). Apply these rules and solve for the unknown element.

An oxidising agent (oxidant) accepts electrons — it gets reduced (its oxidation state decreases). A reducing agent (reductant) donates electrons — it gets oxidised (its oxidation state increases). A useful memory trick: OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

A disproportionation reaction is one in which the same element is simultaneously oxidised and reduced. For example, in 2H₂O₂ → 2H₂O + O₂, oxygen goes from −1 (in H₂O₂) to −2 (in H₂O) and also to 0 (in O₂). So oxygen is both oxidised and reduced.

Split the reaction into two half-reactions: one for oxidation and one for reduction. Balance each half-reaction separately (balance atoms, then add H₂O for O, H⁺ for H in acidic media, or OH⁻/H₂O in basic media, then add electrons to balance charge). Multiply the half-reactions to equalise electrons, then add them together and cancel common terms.

The electrochemical series (electromotive series) is a list of elements (and their ions) arranged in order of their standard reduction potential (E°). Elements higher in the series (more positive E°) are better oxidising agents. Elements lower in the series (more negative E°) are better reducing agents. A metal can displace another metal from its salt solution only if it is above the other metal in the activity series (lower reduction potential).

A galvanic (voltaic) cell converts chemical energy into electrical energy spontaneously (ΔG < 0). Oxidation occurs at the anode (negative terminal in galvanic cells) and reduction at the cathode (positive terminal). An electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction (ΔG > 0). In electrolytic cells, the anode is connected to the positive terminal of the battery.