Chemical Equilibrium
This section covers Chemical Equilibrium: the law of mass action, Kc and Kp, reaction quotient, ICE tables, degree of dissociation and Le Chatelier's principle. Numerical questions often combine equilibrium amounts with pressure or volume changes.
Core rules
- For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d/([A]^a[B]^b). Pure solids and pure liquids are omitted.
- Kp = Kc(RT)^Δn, where Δn is the change in gaseous moles.
- Q compared with K predicts direction: Q < K shifts forward; Q > K shifts backward; Q = K means equilibrium.
- Le Chatelier: changes in concentration, pressure, volume or temperature shift the equilibrium. A catalyst changes only the rate of approach, not K.
Worked example
PCl₅(g) ⇌ PCl₃(g) + Cl₂(g). Start with 1.00 mol PCl₅ in a 1.00 L vessel. At equilibrium 40% has dissociated: PCl₅ = 0.60 mol, PCl₃ = 0.40 mol, Cl₂ = 0.40 mol. Then Kc = (0.40 × 0.40)/0.60 = 0.267. Since Δn = 1, Kp = KcRT. If the vessel is enlarged, Q < K and the equilibrium shifts toward the side with more gas moles, increasing dissociation.
Traps
Do not include solids or pure liquids in K; do not use initial concentrations in K; do not assume pressure changes shift equilibria when Δn = 0; remember that temperature, not pressure, changes the value of K.