Chapter 7: Redox Reactions | Summary

A more modern and precise way to define redox reactions is through electron transfer. Oxidation is the loss of electrons by any species, and reduction is the gain of electrons by any species. An oxidising agent is an acceptor of electrons, while a reducing agent is a donor of electrons. Each redox reaction can be split into two half-reactions — one involving oxidation (loss of electrons) and the other involving reduction (gain of electrons).

Competitive electron transfer reactions illustrate the relative tendency of metals to lose electrons. For example, zinc displaces copper from copper sulphate solution because zinc has a greater tendency to lose electrons than copper. Similarly, copper displaces silver from silver nitrate solution. This leads to the concept of the electrochemical series, which ranks metals in order of their electron-releasing tendency: Zn > Cu > Ag.

Oxidation number (also called oxidation state) is a concept used to track electron shifts in chemical reactions, especially those involving covalent compounds where complete electron transfer does not occur. It is defined as the charge an atom would have if the shared electron pair were assigned entirely to the more electronegative atom. A consistent set of rules is used to assign oxidation numbers: free elements have oxidation number zero; monoatomic ions have oxidation number equal to their charge; oxygen is usually -2 (except in peroxides where it is -1, superoxides -1/2, and in OF2 it is +2); hydrogen is +1 except in metal hydrides where it is -1; fluorine is always -1; the algebraic sum of oxidation numbers in a neutral compound is zero, and in a polyatomic ion equals the charge on the ion.

Using oxidation numbers, oxidation is defined as an increase in oxidation number, and reduction as a decrease. Stock notation, introduced by Alfred Stock, expresses the oxidation state of a metal using Roman numerals in parentheses after the metal symbol — for example, Fe(II)O and Fe(III)2O3. Fractional oxidation states, such as in C3O2 (4/3 for carbon) or Fe3O4 (8/3 for iron), represent average values and actually indicate the presence of the element in two or more different integer oxidation states within the compound.

Redox reactions are classified into four types. Combination reactions involve two or more substances combining to form a single product, with at least one reactant in elemental form — for example, C + O2 → CO2. Decomposition reactions are the reverse: a compound breaks down into simpler substances, at least one of which is elemental — for example, 2KClO3 → 2KCl + 3O2. Not all decomposition reactions are redox reactions; decomposition of CaCO3 is not a redox reaction as there is no change in oxidation numbers.

Displacement reactions involve an atom or ion in a compound being replaced by another. Metal displacement reactions (e.g., Zn displacing Cu from CuSO4) are widely used in metallurgy. Non-metal displacement includes hydrogen displacement from water or acids by active metals, and halogen displacement where a more reactive halogen displaces a less reactive one from its salt — for example, Cl2 displacing Br- and I- ions, forming the basis of the 'Layer Test'. Disproportionation reactions are special redox reactions in which a single element in one oxidation state is simultaneously oxidised and reduced — for example, 2H2O2 → 2H2O + O2, where oxygen in the -1 state goes to both -2 and 0 states.

Two methods are used to balance redox equations. The Oxidation Number Method involves writing correct formulas, identifying atoms that change oxidation number, calculating increases and decreases, multiplying to equalise changes, balancing ionic charges by adding H+ (acidic medium) or OH- (basic medium), and finally balancing hydrogen atoms by adding water molecules. The Half-Reaction (Ion-Electron) Method splits the equation into two half-reactions — one for oxidation and one for reduction — each of which is balanced separately for atoms and then for charges by adding electrons. The half-reactions are multiplied to equalise electrons and then added.

For reactions in acidic medium, H+ ions and H2O are used to balance the half-reactions. For basic medium, the equation is first balanced as in acidic conditions, and then for each H+ ion present, an equal number of OH- ions are added to both sides; where H+ and OH- appear together on the same side, they are combined to give H2O. Both methods give the same final balanced equation.

Redox titrations are used to determine the strength of an oxidant or reductant using a redox-sensitive indicator. Permanganate ion (MnO4-) acts as a self-indicator since it is intensely coloured and its pink colour disappears when fully reduced. Dichromate (Cr2O72-) uses an external indicator such as diphenylamine. A third method uses iodometric titration, where Cu2+ oxidises I- to I2, which is then titrated against thiosulphate (S2O32-) using starch as indicator.

The concept of electrode processes arises when a redox reaction is conducted in a way that prevents direct electron transfer. In the Daniell cell, zinc and copper rods are placed in their respective salt solutions connected by a salt bridge. Electrons flow from the anode (zinc, where oxidation occurs) through the external circuit to the cathode (copper, where reduction occurs). The potential difference between electrodes is the electrode potential. Standard electrode potential (E°) is measured at 298 K with all species at unit concentration and 1 atm for gases; H+/H2 electrode has E° = 0.00 V by convention. A negative E° indicates a stronger reducing agent than H+/H2; a positive E° indicates a weaker reducing agent.

Redox reactions are a fundamental class of chemical reactions characterised by the simultaneous occurrence of oxidation and reduction. The concept has evolved through three tiers: classical (based on oxygen and hydrogen transfer), electronic (based on electron transfer), and oxidation number (tracking electron shifts in covalent systems). Redox reactions are classified into combination, decomposition, displacement, and disproportionation types, and can be balanced using either the oxidation number method or the half-reaction (ion-electron) method.

From a board examination perspective, this chapter is highly important as it provides the foundation for electrochemistry studied in Class XII. Students must master the rules for assigning oxidation numbers, identifying oxidants and reductants, classifying redox reactions, and balancing equations in both acidic and basic media. Understanding the Daniell cell and standard electrode potentials is also essential for predicting the feasibility of redox reactions and for practical applications like batteries, corrosion, and industrial electrochemical processes.