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In a general chemistry lab protocol (green chemistry, ISBN-13: 978-1482230208), students determine the content of ascorbic acid in a vitamin C tablet. There is an ascorbic acid standard in a buret, and there is also some ferric ammonium sulfate solution at known concentration, i.e. iron(III). The indicator is salicylic acid, which forms a colored complex with iron(III), but not iron(II).

The suggested protocol is a back titration where you first add a known excess amount of ferric salt to the unknown amount of ascorbic acid and add the indicator, giving a colored solution. Then, you titrate back with the ascorbic acid standard. The endpoint is the disappearance of color.

Is there are reason for the back titration in this case? In other back titrations, there is a step after adding an excess of reactant (for calcium carbonate, you add an excess of HCl and then drive out the carbon dioxide with heat, making the back titration a strong acid/strong base reaction). For the ascorbic acid reacting with ferric salt, I wonder if the endpoint is easier to see, or maybe the kinetics of complex formation motivate this (it says to let the solution sit before starting the back titration).

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    $\begingroup$ In my opinion, there is a mistake in the type of iron ions reacting with salicylic acid. Salicylic acid is an indicator that makes a red complex with iron(III) and not with iron(II), as stated in the first section. $\endgroup$
    – Maurice
    Commented Aug 4, 2022 at 19:41
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    $\begingroup$ You are right Maurice, the text "Green Chemistry" clearly states, iron(III)-salicylate acid complex along with its structure. $\endgroup$
    – AChem
    Commented Aug 4, 2022 at 19:48
  • $\begingroup$ @Maurice I'll fix that. $\endgroup$
    – Karsten
    Commented Aug 4, 2022 at 20:11

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The idea of using green chemistry to titrate vitamin C with environmentally friendly reagents is understandable, but it may not be the most appropriate method. Iodometry has traditionally been the traditional method for ascorbic acid analysis. The main reason for doing back-titration is that ferric ion is a very gentle oxidizing agent but good enough for ascorbic acid. Therefore, we wish to ensure that all of the acid in the sample is oxidized by ferric ions.

Now note that this is a pseudo-back titration, because the titrant and the analytes are the same! In other words, after adding excess iron (III) to the sample solution, a "direct" titration is being conducted between ascorbic acid and iron (III). I assume the titrant concentration is higher than ferric ammonium sulfate to speed up the reaction and ensure completion of the redox process. As an exercise, why don't you ask the students to try direct titration and back-titration and then do a statistical test to see if there is a difference. I mean half the class perform direct titration and half the class do this pseudo-back titration.

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  • $\begingroup$ Excellent idea! Is there a preference for titrations going from colorless to color or vice versa? $\endgroup$
    – Karsten
    Commented Aug 4, 2022 at 20:15
  • $\begingroup$ No I don't think so, the color change must be rapid, that is it. $\endgroup$
    – AChem
    Commented Aug 4, 2022 at 20:20
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    $\begingroup$ I changed the title to pseudo-back titration. $\endgroup$
    – Karsten
    Commented Aug 4, 2022 at 20:42
  • $\begingroup$ @Karsten-apprentice, Did you get different results of back vs. direct titrations? $\endgroup$
    – AChem
    Commented Aug 24, 2022 at 14:24
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    $\begingroup$ You are giving me a lot of J, Chem. Ed. suggestions :-). I'm happy with the one I already published, but maybe another one some day. Another comment I will have to delete, I should stop writing them. $\endgroup$
    – Karsten
    Commented Aug 24, 2022 at 15:30

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