Question

The half-equivalence point of a titration occurs half way to the end point, where half of the analyte has reacted to form its conjugate, and the other half still remains unreacted. If 0.200 moles of a monoprotic weak acid (Ka = 5.3 × 10-5) is titrated with NaOH, what is the pH of the solution at the half-equivalence point?

Answer 1

Use the Henderson-Hasselbach equation
pH = -log(Ka)+log(A/HA)

Here, [HA] is theHere, [HA] is the molar concentration of the undissociated weak acid, [A⁻] is the molar concentration (molarity, M) of this acid's conjugate base and pK_a is −log₁₀K_a where K_a is the acid dissociation constant

In this case fraction of the molar concentration of the undissociated weak acid and the molar concentration (molarity, M) of this acid's conjugate base   [A/HA] is 1, so its just pH = -log(Ka)

pH = -log[5.3x10^-5]

pH = 4.27