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If I understand this titration correctly, if we titrated 50mL 1.00M acetic acid with 1.00M NaOH, the equivalence point would happen at a pH greater than 7, and less than 50mL of NaOH would need to be used.

What I'm curious about is why the titration does not take 50mL of both and why it does not end at pH of 7. I can accept the numbers above, but I would like to know why.

The reason I struggle with this is because since the sodium hydroxide is a strong base and will ionize almost entirely, its hydroxide ions will react with whatever hydrogen ions have been ionized from the acetic acid. This neutralization will force more acetic acid to be ionized, and the process will repeat. Then, the same volume and molarity of each will be used, and the titration will end with complete neutralization at pH 7. Why is this not the case? Can weak acids and bases not be forced to ionize in such a way? Is there some point at which a weak base or acid will not ionize any more, no matter what?

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marked as duplicate by airhuff, Ben Norris, ron, Klaus-Dieter Warzecha, M.A.R. Apr 25 '17 at 13:57

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    $\begingroup$ Related:chemistry.stackexchange.com/questions/47705/… $\endgroup$ – Tyberius Apr 24 '17 at 14:53
  • $\begingroup$ @Tyberius Thanks for that. It definitely is related, although it is not the same question I'm asking. I had not thought of the buffer effect in there, but it does make sense, yes. $\endgroup$ – Jesuspowder Apr 24 '17 at 15:02

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