Chemical equilibrium occurs when the rates of forward and reverse reactions are equal, resulting in no net change in the concentrations of reactants and products.

The equilibrium constant (K) quantifies the ratio of product to reactant concentrations at equilibrium. Its value indicates the extent to which a reaction proceeds.

The reaction quotient (Q) is calculated similarly to K but uses current (not equilibrium) concentrations. Comparing Q to K predicts the direction of reaction:

• If Q < K, the reaction proceeds forward (toward products).

• If Q > K, the reaction proceeds in reverse (toward reactants).

The standard Gibbs free energy change (ΔG;) relates to K by:

• ΔG;=ΔG;⁰+RTlnQ

• At equilibrium: ΔG;=0 and Q=K

• ΔG;⁰=−RTlnK

K > 1: products favored at equilibrium; K < 1: reactants favored.

Le Châtelier’s Principle: If a system at equilibrium is disturbed, it will shift to counteract the disturbance and restore equilibrium.

Adding reactants or products affects Q and shifts the equilibrium:

• Adding reactant: Q decreases, reaction shifts forward to produce more product.

• Adding product: Q increases, reaction shifts backward to produce more reactant.

For reactions involving gases, equilibrium expressions use partial pressures; for solutions, concentrations in molarity are used.

For a general reaction aA+bB⇌cC+dD:

• In solution: K=[C]c[D]d[A]a[B]b

• In gases: K=P_CcP_DdP_AaP_Bb

When combining reactions, the equilibrium constant for the overall reaction is the product of the equilibrium constants for the individual steps.

The magnitude of K is related exponentially to the standard free energy change: K=e−ΔG;⁰RT