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During a titration lab where 0.1 N of NaOH is added to 20 mL of acetic acid with the same concentration, the equivalence point occurred after adding about 70ish milliliters of NaOH (different peers acquired different values, ranging from 65 mL and 75 mL of NaOH). This completely puzzled me, because after calculating a theoretical titration curve, my equivalence point occurred after adding about 20 mL of NaOH.

My teacher could not explain why this happened, except that acetic acid can act as a buffer – but how do you calculate this and incorporate it into a theoretical titration curve? I just don't understand how 20 mL of acetic acid would require 70 mL of NaOH to neutralise it (both have the same concentration).

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    $\begingroup$ There is nothing about acetic acid that would make your titre value over 3 times the expected value. I would be willing to bet that someone messed up somewhere with the chemicals you were given. $\endgroup$ – orthocresol Nov 28 '15 at 23:51
  • $\begingroup$ That is the only explanation I can come up with as well. It is just unfortunate that I will loose marks for not having a titration curve that does not reflect the curve acquired during the lab. I will try and convince my prof that the NaOH given to me must have been diluted incorrectly. $\endgroup$ – Denis Nov 28 '15 at 23:59
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It can't happen. Neutrality in aqueous solution means equal quantities H+ and OH-; since your solution is unlikely to have a net electric charge, so too must your anion and cation quantities be equal. Buffering from acetic acid being a weak acid is meaningless if you don't have other anions around to make up that balance.

Solution is electrically neutral: [H+] + [Na+] = [OH-] + [AcO-]

Solution is pH neutral: [H+] = [OH-]

Substitute the latter in the former, subtract, you can only get: [Na+] = [AcO-]

If the concentrations of both solutions were the same, and you're titrating by adding both to the same volume, you can only have equal concentrations if you add equal volumes.

If this was not the case (as you describe) we have to violate one of our assumptions:

  • Solution actually had a net electric charge (you're doing the titration in space and you're bombarded by beta radiation, i.e. unlikely)
  • Solution was not truly neutral. Maybe someone calibrated your pH meters wrong, maybe everyone titrated with a faucet while texting and over-shot by a factor of 3.
  • Contamination with a strong acid added a variable we don't know about to the ion balance.
  • Most likely, the solutions were not truly the same concentration. See if your teacher made her NaOH solution from solid NaOH that she let sit around in open air. It can pick up a lot of moisture from the air, i.e. it's hygroscopic, and before you know it the weight you're measuring on a balance is half water.
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  • $\begingroup$ Thank you! I agree with everything you said. I was looking at the data acquired from a titration using the same NaOH with HCl, and I am beginning to think that the acetic acid may have been diluted incorrectly... but then again, I remember doing a previous lab with the NaOH and I acquired weird results... so the NaOH may very well have absorbed moisture from somewhere (we use the pellets) $\endgroup$ – Denis Nov 29 '15 at 17:41

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