Question

Regarding buffers...

Answer 1

I'm a bit confused about buffers. I realize that the purpose of a buffer is to resist changes in pH by neutralizing other species. However, what I do not understand is this. Given a weak acid, HA...\[HA(aq)+H_2O(l) \rightarrow H_3O^+(aq)+A^-(aq)\]Given this equation, we know that the concentrations of each aqueous species can be accommodated for with...\[K_a(HA)=\frac{[H^+][A^-]}{[HA]}\]Where the concentrations are at equilibrium.

This is where my understanding breaks down. The next step in making a buffer solution is to add some species which contains the conjugate base of the weak acid (in this case, A^-). For instance...
\[BA(s) \rightarrow B^+(aq)+A^-(aq)\]We'll make the assumption here that B+ is the conjugate acid of a very strong base (100% association). In this case, we are adding A- to the solution and thus forcing the first reaction...\[HA(aq)+H_2O(l) \rightarrow H_3O+(aq)+A^-(aq)\] to shift to the left, according to Le Châteliers principle. In this case, the reaction is always going to be heavily shifted to the left because the K_a of HA is low.

I know that the hydronium concentration comes into play here, but I'm unsure upon how it changes exactly to accomodate and create a buffer solution. Could anyone clear this up for me?

Answer 2

I guess my main question is regarding how it's possible to get equal amounts of weak acid and conjugate base? The concentration of hydronium would have to increase significantly (and the pH decrease) for this to occur, would it not?