An aqueous solution initially contains 0.2 M of a weak dibasic acid H₂A . A strong acid is added to the solution without significantly changing its volume until the equilibrium concentration of the fully deprotonated anion A²⁻ reaches exactly 4.0 10⁻¹² M. If the first and second ionization constants of H₂A are 4.0 10⁻⁶ and 5.0 10⁻¹⁰ respectively, the pH of the resulting solution is:

Step-by-step solution

The overall dissociation reaction for the weak dibasic acid is: H₂A 2H^+ + A²⁻ The overall equilibrium constant K_ net is: K_ net = K_ a1 K_ a2 = (4.0 10⁻⁶) (5.0 10⁻¹⁰) = 20.0 10⁻¹⁶ = 2.0 10⁻¹⁵ The equilibrium expression is: K_ net = [H^+]^2[A²⁻] [H₂A] Due to the presence of the strong acid, the dissociation of H₂A is highly suppressed. Therefore, the equilibrium concentration of H₂A is approximately equal to its initial concentration: [H₂A] 0.2 M We are given the equilibrium concentration of the anion: [A²⁻] = 4.0