What is the pH half way to, and at the stoichiometric point? no idea how to do this problem.?

Question

A.) 101.0 mL of a solution of quinine contains 202.0 x 10-3 mol of quinine.
If this is titrated with 0.200 M HCl, what is the pH half-way to the stoichiometric point ?

Quinine is monobasic with Kb = 7.90 x 10-6

B.) What is the pH at the stoichiometric point ?

Answer 1

What is it with quinine? About half a dozen folks have asked this question or somthing very similar.

Quinine is an organic substance. If we call it Q,
and dissolve it in water, this reaction occurs:
Q + H+ -> QH+

to a slight extent. Thus it raises the OH- conc, lowers the H+ conc, since [H+][OH-]= 10x10-15 in water.

Addition of a strong acid converts more Q to QH+. The stoichiometric point occurs when all the Q is so converted. At this point, the ability of the solution to resist change of pH due to acid is used up.

The equation with Kb is
[QH+][OH-]/[Q] = Kb

At the halfway point, [QH+]=[Q]. Then [OH-] = Kb
From the water ionization, [H+]= 1.3 x10-9 (appx)
and pH= 9 - log 1.3 = 8.85 appx.

At the stoichiometric point, pH=7. However, the next "chunk" of 0.2 M HCl will change the pH notably