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I have been preparing lessons for some of my students about preparing acidic and basic solutions, and I keep finding pre-made labs in which the students "standardize a basic solution" by titrating it with an acid of "known" concentration.

The problem is that the students create a solution of NaOH by measuring a certain mass and adding it to a volumetric flask, filling with water to the desired volume. But the students will do the same thing with an acid (potassium hydrogen pthalate, for example). They will measure out a specific mass, do the calculations to determine the molarity of the KHP solution, and then do the titration to determine the molarity of the NaOH solution. But if the students took the mass of the NaOH in the beginning, can they not just use that information to determine the approximate molarity of the solution?

As the title says, why bother to standardize? If our NaOH solution had some room for error in measurement and calculation, surely titrating against another solution of "known" concentration would only introduce more error, no?

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Usually, if we're doing these for basic application or experimentation, standardization doesn't really matter. But when it comes to anything analytical where you start to involve calculations, standardization is a must. This is done with NaOH because it's hygroscopic and readily sucks up the moisture in the air. So what is being weighed isn't totally NaOH, but also the moisture that it has absorbed. So almost always the concentration will be lesser than what is sought to be prepared because of this. Tedious as it may be (because you also have to heat up the KHP for around an hour before titrating), standardization brings you closer to the true concentration, but not exactly on the mark. At least in the end, calculations using the standardized concentrations will be analytical. This is why they never introduce standardization in general chemistry courses, I only got to grips with it when I took up Analytical Chemistry, and back then it was really tedious to do.

For example, I actually experienced this before where we had to prepare a 1M stock of NaOH, but after standardization it turned out to be around 0.8 (maybe due to the humidity in the lab at the time).

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  • $\begingroup$ Fascinating take on it. My textbook is a high school textbook from about 2008, I think. Almost every lab the publishers provided with the book involves standardization, which left me awfully curious, as I had only done it once before myself. $\endgroup$ – Jesuspowder Apr 23 '17 at 2:16
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    $\begingroup$ I think it'd be something great to teach your students. Then you could follow it up with another experiment that uses the standardized solutions like the titration of Vitamin C for example if they're at that level. $\endgroup$ – tumblewush Apr 23 '17 at 2:22
  • $\begingroup$ I don't doubt they're at that level, but my budget isn't at that level. Maybe next year. Thanks for the great idea! $\endgroup$ – Jesuspowder Apr 23 '17 at 2:25
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Nope, it will not introduce, it will decrease the error.

There are a list of compounds which are used for preparation of various standard solutions and appropriate protocols. The chemical structure for such compounds does not depends on external factors (following the protocol) and can be used to determine the concentration precisely. Additionally more sensitive methods can be used to determine the final concentration of standard solutions, e.g., UV-vis spectroscopy.

To understand why you don't want to calculate the concentration of sodium hydroxide solution based on its weight and why this method will introduce more error, just let solid NaOH sit on a bench top for a little bit (please use a weighting paper/dish for it). Let me know what did you find out! (:

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  • $\begingroup$ I'm guessing it's hydroscopic, and that might introduce some mass error. I had not thought of that. Since my NaOH is from the late 90s, it might be wise to do this after all... $\endgroup$ – Jesuspowder Apr 23 '17 at 2:13

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