Chemical equilibrium: study guide

Key ideas

At equilibrium the forward and reverse reactions go at the same rate and the concentrations stop changing. The equilibrium constant K tells how far a reaction goes, and Le Chatelier’s principle predicts how it shifts.

• For aA + bB ⇌ cC + dD, K_c = [C]ᶜ[D]ᵈ ÷ ([A]ᵃ[B]ᵇ). Pure solids and liquids are left out.
• Compare the reaction quotient Q with K. If Q < K the reaction goes forward, if Q > K it goes backward, and if Q = K it is at equilibrium.
• Le Chatelier: a system under stress shifts to relieve it. Adding reactant shifts the equilibrium forward, and raising the temperature favors the endothermic direction.
• For gases K_p = K_c(RT)^Δn, and for a sparingly soluble salt K_sp is the product of the ion concentrations raised to their coefficients.

Key equations

• K_c = [C]ᶜ[D]ᵈ ÷ ([A]ᵃ[B]ᵇ)
• Q vs K
• K_p = K_c(RT)^Δn
• K_sp = [M]ᵐ[X]ˣ

Worked example

For N₂ + 3 H₂ ⇌ 2 NH₃, the equilibrium concentrations are [N₂] = 0.50, [H₂] = 0.20 and [NH₃] = 0.10 mol/L. K_c = 0.10² ÷ (0.50 × 0.20³) = 0.010 ÷ 0.0040 = 2.5. For AgCl with K_sp = 1.8 × 10⁻¹⁰, the solubility is s = √K_sp = 1.3 × 10⁻⁵ mol/L.

Common mistakes

• Putting reactant concentrations in the numerator instead of the products.
• Including pure solids or liquids in the expression for K.
• Thinking a catalyst changes the value of K. It speeds up both directions equally and only helps equilibrium be reached sooner.

Practice

Test yourself with practice questions on this topic: Chemical equilibrium