Redox reactions form an important class of reactions in which oxidation and reduction occur simultaneously. The chapter presents three tiers of conceptualisation:
- Classical idea: Oxidation is the addition of oxygen/electronegative element to a substance, or the removal of hydrogen/electropositive element from it. Reduction is the reverse. Because both always occur together, the term 'redox' was coined.
- Electron-transfer idea: Oxidation is the loss of electron(s) by a species; reduction is the gain of electron(s). An oxidising agent (oxidant) is an electron acceptor; a reducing agent (reductant) is an electron donor. Each step can be written as a half reaction. Competitive electron transfer experiments (Zn in CuNO3, Cu in AgNO3) give the reducing order Zn > Cu > Ag.
- Oxidation number idea: Oxidation number is assigned by a consistent set of rules (free element = 0; monatomic ion = its charge; H = +1 except -1 in metal hydrides; O = -2 except -1 in peroxides, -1/2 in superoxides, positive with F; F always -1; sum = 0 for a neutral compound, = charge for an ion). Oxidation is an increase in oxidation number; reduction is a decrease. An oxidant increases, a reductant decreases, the oxidation number of an element.
Redox reactions are classified into four types: combination, decomposition, displacement (metal and non-metal) and disproportionation. Not all combination/decomposition reactions are redox.
Redox equations are balanced by the oxidation number method and the half reaction (ion-electron) method. Redox reactions are the basis of redox titrations (MnO4- self indicator, diphenylamine for Cr2O7^2-, starch for iodine). The concepts of redox couple, electrode and standard electrode potential (H+/H2 = 0.00 V) introduce electrode processes and cells such as the Daniell cell, studied further in Class XII.