Chemical equilibrium is a dynamic process where the rates of the forward and reverse reactions are equal, resulting in no net change in concentrations of reactants and products.

Equilibrium can be homogeneous (all species in the same phase) or heterogeneous (species in different phases).

The equilibrium constant (K) expresses the ratio of product concentrations to reactant concentrations, each raised to the power of their coefficients in the balanced equation. Solids and pure liquids are omitted from K expressions.

K=[C][D][A][B] for a reaction aA + bB ⇌ cC + dD.

K is temperature dependent; changing temperature alters the value of K.

The magnitude of K indicates whether products or reactants are favored at equilibrium: K >> 1 favors products, K

• Kp is used for gases (partial pressures), Kc for concentrations in molarity. They are related by: Kp=Kc(RT)(Δn) where Δn = moles of gaseous products - moles of gaseous reactants.

• Le Châtelier's Principle: If a system at equilibrium is disturbed (by changes in concentration, pressure, or temperature), the system shifts to counteract the disturbance and re-establish equilibrium.

• Adding or removing reactants/products, changing pressure/volume, or temperature can shift the equilibrium position.

• For exothermic reactions, increasing temperature shifts equilibrium toward reactants; for endothermic, toward products.

• Equilibrium calculations may involve determining K from concentrations/pressures, or finding unknown concentrations/pressures given K and other values.

• When combining multiple reactions, the overall K is the product of the K values for each step, adjusted for any reversal or multiplication of reactions.

• Practice problems often require writing K expressions, predicting shifts in equilibrium, and performing calculations using the relationships above.