When adding a strong base to a buffer solution, with a generic weak acid HA and generic weak conjugate base A-, is the initial amount of H+ ignored because it is so small? Specifically, at least in my text book, we only consider what happens to HA and A- when a strong base, say NaOH, is added.
Suppose you have a weak acid and its corresponding base in solution. Suppose the solution has the pH = 5.
HA + H2O = A- + H3O+
Now we add a strong base to the solution. What will happen?
The pH will increase, suppose it will increase to pH = 6.
The weak acid-base system will immediately respond to the pH change of the solution and a new equlibrium will be established defined by the new pH and the equilibrium constant of the weak acid-base system. The equilibrium will be forced to the right, i.e. more A- will be fomed at pH = 6 compared to what we had at pH = 5.
Ok, what has happened to the hydrogen ions present at pH = 5 or those that were formed up to
pH = 6?
They have reacted with the strong base forming water.
Do we think the concentration of the hydroxide ions from the strong base has decreased at pH = 6?
At pH = 6 the hydrogen ion concentration is lower than at pH = 5. This is equal to say that the hydroxide ion concentration is higher at pH = 6 than at pH = 5. The difference in the amount of hydroxide ions present at pH 6 and 5 will be equal to the amount of hydroxide ions consumed in the transfer from pH 5 to 6. What you put in, you get back in this case. The net effect is that the hydroxide ion concentration of the strong base has not changed at the equilibrium of the process.
I suggest you also look at my answer for another question "Buffer and strong acid".