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Many chemical reactions are reversible — they proceed in both forward and reverse directions simultaneously until the net rates equalize at equilibrium. The equilibrium constant K is the ratio of product concentrations to reactant concentrations, each raised to their stoichiometric coefficients. Le Chatelier's principle states that a system at equilibrium shifts to counteract any applied stress (change in concentration, pressure, or temperature). The reaction quotient Q, compared to K, indicates whether a system will shift forward (Q < K), backward (Q > K), or is already at equilibrium (Q = K).
Practice ICE (Initial-Change-Equilibrium) tables to solve for equilibrium concentrations. Apply Le Chatelier's principle qualitatively to predict shifts for various stresses. Use the small x approximation when Ka is very small relative to initial concentration, but verify its validity.
Chemical equilibrium is one of the most conceptually rich ideas in general chemistry because it forces you to think about reactions as ongoing, two-directional processes rather than one-way events. When you mix nitrogen and hydrogen gas at high temperature, both the forward reaction (making ammonia) and the reverse reaction (decomposing ammonia) happen simultaneously. Equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction — not when the reaction "stops."
The equilibrium constant K captures the outcome of this balance. For a reaction aA + bB ⇌ cC + dD, the equilibrium expression is K = [C]c[D]d / [A]a[B]b, where the brackets denote molar concentrations at equilibrium and the exponents are the stoichiometric coefficients. K is a fixed number at a given temperature — large K means products predominate at equilibrium, small K means reactants predominate. Notice that K depends only on temperature; changing concentrations or pressure shifts where equilibrium sits but does not change the value of K.
Le Chatelier's principle is the conceptual shortcut: any stress applied to a system at equilibrium will be "counteracted" by a shift in the equilibrium position. If you add reactant, the system shifts forward. If you remove product, the system shifts forward. If you increase pressure (in a gas-phase reaction), the system shifts toward the side with fewer moles of gas. The mathematical reason this works is the reaction quotient Q. At any moment, Q = [products]/[reactants] using current (not equilibrium) concentrations. If Q < K, the system shifts forward; if Q > K, it shifts in reverse; if Q = K, it is at equilibrium.
ICE tables give you a systematic way to calculate equilibrium concentrations. Set up rows for Initial concentration, Change in concentration (−x for reactants, +x for products, scaled by stoichiometry), and Equilibrium concentration. Substitute the equilibrium row into the K expression and solve for x. When K is very small (≤ 10⁻⁴), the "small x approximation" lets you drop x from sums and differences, simplifying the algebra — but always check that x is indeed small relative to the initial concentrations after solving.
One crucial distinction: temperature is unique among the stresses you can apply. Adding more reactant, changing pressure, or introducing an inert gas shifts the position of equilibrium but leaves K unchanged. Changing temperature actually changes the value of K — it alters the equilibrium constant itself. Whether K increases or decreases with temperature depends on whether the forward reaction is endothermic or exothermic, which connects this topic directly to thermodynamics.