Chemical Equilibrium || Reaction quotient and Extent of Reaction || Lecture # 1 || Dr Rizwana

Chemical equilibrium isn’t a “done” state where reactions stop—it’s a dynamic balance where forward and reverse reactions continue at equal rates, leaving reactant and product concentrations unchanged at a specific temperature and pressure. The lecture frames equilibrium as a condition that appears in reversible reactions, including cases where reactants never fully convert because products decompose back into reactants. A classic illustration is the decomposition of PCl5 into PCl3 and Cl2, represented with a double-headed arrow to show both forward and backward directions operating simultaneously.

As the reaction proceeds, reactant concentration falls while product concentration rises until a point is reached where neither changes further. At that stage, the system is said to have achieved equilibrium: the concentrations remain constant even though molecules are still reacting. This is labeled “dynamic equilibrium,” emphasizing that the forward rate equals the reverse rate. The lecture also distinguishes equilibrium from the idea of a fixed 50–50 mixture; equilibrium composition depends on conditions and can favor either reactants or products.

To describe equilibrium quantitatively, the lecture introduces the equilibrium constant, Kc, defined as the ratio of product concentrations (each raised to its stoichiometric power) to reactant concentrations (also raised to stoichiometric powers). Kc is tied to the law of mass action, and the lecture uses the formation of hydrogen iodide from hydrogen and iodine as an example, showing how stoichiometric coefficients determine the exponents in the Kc expression. Graphically, the reactant concentration curve decreases toward a constant value, while the product curve increases toward a constant value; those plateaus correspond to equilibrium.

The magnitude of Kc becomes a predictor of equilibrium composition. When Kc is very large (for example, greater than 10^3), products dominate at equilibrium, meaning the reaction proceeds far toward completion. When Kc is very small (for example, less than 10^-3), reactants dominate because the reverse direction is favored. When Kc falls between these extremes (roughly between 10^-3 and 10^3), both reactants and products can coexist in significant amounts at equilibrium. The lecture gives specific comparisons using Kc values for reactions forming water, hydrogen chloride, hydrogen bromide, and hydrogen iodide, and contrasts them with a reaction like formation of nitrous oxide where Kc is extremely small.

Finally, the lecture connects equilibrium constants to reaction direction using the reaction quotient, Qc. Qc has the same mathematical form as Kc, but it can be calculated at any time, not only at equilibrium. If Qc is less than Kc, the system has more “room” to form products, so the reaction proceeds forward. If Qc is greater than Kc, the system shifts in the reverse direction to reduce Qc. When Qc equals Kc, the system is at equilibrium for that specific time and set of conditions. This relationship allows prediction of whether a reaction mixture will move toward more products or more reactants at a chosen moment—an essential tool for understanding the extent of reaction and the direction of change under equilibrium conditions.