Question #b36f6

The image below shows a typical titration curve for the addition of a strong base into a weak acid. Note that the equivalence point comes at a value somewhat greater than seven.

Two points on the graph are of interest. The first is called the half-equivalence point. (Sorry, it is not shown on the graph, but it lies midway from the pH axis to the equivalence point.) This is reached when the amount of base added in enough to change half of the acid into its conjugate base. At this point #[A^-] = [HA]#, and the pH value that is equal of the log of the acid's #K_a# value .
(Check the #K_a# expression to verify this.)

If one acid chosen is weaker than another, its #K_a# value is smaller, and the log of this value is larger. This means the half-equivalence point comes at a higher pH for a weaker acid, and the rather flat portion of the graph is at a higher pH .

The second point of interest is, of course, the equivalence point. This comes at a point where the amount of added base matches the original amount of acid present. At this point, essentially all the HA has been ionized into #A^_# and the solution is identical to one that was prepared using the weak base #A^-#.

So, the calculation of the pH at the equivalence point is identical to the calculation of pH of a solution made by dissolving #A^-# in water. A weaker acid will have a stronger conjugate, and so, the pH of the equivalence point rises if the chosen acid is weaker.

So, to summarize, as the #K_a# value gets smaller for a weaker acid, expect the pH plot to shift toward higher values of pH, with the flat portion of the graph (where little pH change occurs) coming at a pH equal to the #K_a# value of the acid, and the pH of the equivalence point rising as well.