I'm making a table on excel with the calculated pH of the solution vs the actual measured pH from the lab. My problem is finding the pH of a acetic acid-acetate buffer that has been titrated with HCl, once the moles of HCl are greater than the initial quantity of mols of acetate.
pH = -log(H3O+) = 14-pOH = pKa + log([acetate]/[acetic acid])
Ka = [H3O][acetate]/[acetic acid] = 1.7 x 10^-5
HCl + H20 --> H3O
H3O + CH2COO- = CH2COOH
[acetate] = (initial moles acetate + moles HCl added)/total volume
The Attempt at a Solution
Up to a point I just used the Henderson-Hasselbalch equation. This worked until the moles of HCl were greater than the moles of initial acetate, since this gives a negative number for the concentration of base, which is impossible in reality and mathematically, assuming I was using it correctly. I was getting the concentration of the base by subtracting the initial moles of the base by the moles of HCl and dividing by the total volume.
At that point I turned to the Ka equation. All the HCl goes to H3O, and additional H3O reacts with acetate to form acetic acid as long as there is still acid. So H3O would = [HCl] - [initial acetate]. If that were right, I could take the -log of that for pH. But it doesn't match measured data, or the trend of calculated data up to that point. It gives a more basic result.
I guess the acetic acid is still reacting and adding more H3O, in order to get the Ka. So next I tried:
Ka = ([H3O]+x)(x)/([AA initial]+[initial acetate]-x)
I went through the quadratic with the same sample data I used last time and the x was so small the answer was nearly the same, which is what I would expect, since in this hypothetical the ionization of AA would be supressed by the H3O.
I don't know what I did wrong and am stuck here. I'd really appreciate some help.