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

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

2. The reaction quotient (Q) is calculated similarly to K but uses current 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).

3. The relationship between Gibbs free energy change (ΔG) and equilibrium: ΔG = ΔG^−RTlnQ

   • At equilibrium, ΔG = 0 and Q = K.

   • The magnitude of K is related to the standard free energy change: K = exp(−ΔG^/RT)

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

5. Adding or removing reactants/products affects Q and the direction of the reaction:

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

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

6. Equilibrium expressions for reactions:

   • For a general reaction aA + bB ⇌ cC + dD, the equilibrium constant is: K = [C]c[A]a × [D]d[B]b

   • For reactions in the gas phase, partial pressures are used in the equilibrium expression.

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