[muffins]

Hey guys,
There's an example in my book of pH calculation in a weak acid-strong base titration, and something in there just doesn't really sit well with me at all.
The question is as follows:

Calculate the pH of a solution where 5ml of 0.15M NaOH is titrated onto 25ml of 0.1M HCOOH.  (Ka for HCOOH = 1.8e-4)

The initial pH was calculated in the equilibrium table and after approximations [H₃O⁺]=(Ka[HCOOH])^0.5=4.2e-3M ; -log[H₃O⁺]=2.37 . This stage is clear to me.

Now calculating the moles for OH^- : 0.005l*0.15M=7.5e-4 mol

The reaction: HCOOH(aq)+OH^-(aq) HCO₂^-(aq)+H₂O(l)

Okay here is what I don't quite get... it says - of 0.750 milimoles of OH-, 0.750 milimoles of HCO₂^- are created, and 1.75mmol HCOOH are left.

The book proceeds to calculate the concentrations of the acid and conj. base:
[HCOOH]= 1.75e-3mol/0.03l=0.0583M ; [HCO₂^-]=7.5e-4mol/0.03l=0.0250M
Using HH and after approximations: pH=pKa+log(0.025/0.0583)=3.38

My questions are these:
Why does the book seem to ignore [H₃O⁺] that was initially created just by the weak deprotonation? (the 4.2e-3 M)
Why from 0.75mmol of OH- is 0.75mmol of conj base made? I thought that the initial conj base concentration was equal to the initial [H₃O⁺] and should be accounted for aswell... and I'm not so sure why the stoichiometry here of 0.750mmol OH- = 0.750mmol HCO₂^- is happening.

Does the introduction of OH- neutralize the acid itself? doesn't it neutralize its hydronium ions first? I am so so so confused....

I'd really appreciate some input for this... I've been sitting for hours just trying to get it...

Thank you

[Borek]

We assume neutralization reaction went to completion and final concentrations are defined by the stoichiometry.

It isn't exactly true, but gives reasonably correct results, as long as the acid is not too strong nor too weak, and as long as its concentration is not too low. Say, pKa somewhere between 3 and 11, and concentration of at least 0.001 M.