Final answer:
The increase in H₃O⁺⁻ does not keep pace with the increase in HA because weak acids do not fully dissociate and their conjugate bases can recapture protons. Added strong bases can be consumed by the weak acid, forming its conjugate base and water, which acts as a buffer. The constancy of the ion product of water also minimizes pH changes.
Step-by-step explanation:
The increase in H₃O⁺⁻ does not always keep up with the increase in HA because of the nature of weak acids and their conjugate bases. When HA dissociates in water, it forms H₃O⁺⁻ and A⁺⁻. However, because HA is a weak acid, it does not completely dissociate, meaning not all HA molecules donate their protons to form H₃O⁺⁻. According to Le Chatelier's principle, the reaction is in equilibrium and any increase in H₃O⁺⁻ would shift the reaction to produce more HA, limiting the number of free H₃O⁺⁻ ions. Moreover, the conjugate base A⁺⁻ can recapture protons, creating an equilibrium where there are relatively few free H₃O⁺⁻ ions compared to the amount of dissolved HA.
As a buffer system, when a strong base like MOH is added to a solution containing a weak acid like CH₃COOH, the reaction consumes the base to form the conjugate base CH₃COO⁺⁻ and water. This consumption of the strong base through neutralization limits the increase in pH that would occur if the H₃O⁺⁻ ions were not buffered in this way. The conservation of the ion product of water, which means that [H⁺⁻][OH⁺⁻] must remain constant, also helps regulate the concentration of H₃O⁺⁻ ions, keeping pH changes minimal when a strong acid or base is added to a buffer solution.