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I conducted a titration experiment a few days back and was unable to neutralize the solution in time, so I had to leave the solutions of HCl and NaOH out overnight in order to use them the next day. Also, one of my partners blew into the buret to get a drop of NaOH out. At the end of the experiment, we found that we required 6-7 mL more base to neutralize the HCl solution; another group who was forced to leave their solutions overnight used 3 mL more even than we did.

My hypothesis is that CO2 from the air formed carbonic acid in both solutions, mainly due to the instance where my partner blew into the buret, and thus the acidity of both solutions required that we use more base than we would have originally. As far as I've researched, this seems to be the case. Just to be sure, though: is this a correct assumption? If not, what could be the cause of this discrepancy?

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  • $\begingroup$ @Mithoron Care to explain? Or back your answer with anything? $\endgroup$ – White Fang Jan 12 '17 at 1:31
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    $\begingroup$ Well, sodium alkali draws CO2 from air (carbonates not carbonic acid are created) but it depends on concentration of NaOH and water used, as normally there's CO2 in there from the start. $\endgroup$ – Mithoron Jan 12 '17 at 1:46
  • $\begingroup$ middleschoolchemistry.com/lessonplans/chapter6/lesson10 - How would you explain this, then? Is this property of CO2 not applicable to the solutions present, or what? $\endgroup$ – White Fang Jan 12 '17 at 2:05
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First of all, blowing into the buret seems like a less-than-safe thing to do. $\ce{NaOH}$ could easily fly out into his/her mouth and nose.

I would expect that your acid and base solutions were already at the equilibrium concentration of $\ce{CO_2}$ when you got them, since they were exposed to air. I think that dissolving large amount of $\ce{CO2}$ into solution at atmospheric pressure is no easy task. I just don't think small amounts of $\ce{CO2}$ would account for the difference.

When you say "left them out" do you mean uncovered? If so, there's probably a fair amount of evaporation happening. The analyte solution could be affected depending on your experiment, and the concentration of your titrant would go up, which means you'd actually need less of it to reach the end-point.

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When acquiring both the Acid $\ce{HCl}$ and the base $\ce{NaOH}$, your teacher should have given you a certain amount of either one to start with. Using the data of the given quantity you can acquire the necessary amount of base/acid needed to neutralize the other using the concentration and volume. This is known as the theoretical value needed to carry out the neutralization reaction $\cdots$ and this is probably where from my understanding you misunderstood why you needed more base. Your titration was not wrong as the experimental value (value gotten when experiment is carried out) is usually higher than the theoretical value meaning if 23 ml was required but you used 28 ml, its fine. However the real answer to your question lies in paralax errors (known as reading errors), human errors and friction along the inside of the burret, which causes the discrepancy. Like for example if you wanted to get 200 mL of $\ce{HCl}$ .. you might of gotten 200.5 mL where as the other group got 201 or 202 mL etc.

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