Chemical amounts (A2)
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When you are carrying out an experiment to find the iron content of a solution using potassium manganate(VII), why do you have to boil the vegetables in acid and make this solution up to 100 cm3 with sulphuric acid before performing the titration? Why can’t you just use water? And why, after working out the iron content of the vegetables, is the value obtained way above the daily iron requirement?
Igloo writes
Your “iron solution” obtained after boiling your vegetables with dilute sulphuric acid contains Fe2+ ions.
Unfortunately, these ions are slowly destroyed in solutions with a relatively high pH (pH>7) for two reasons:
First they are readily hydrolysed by water molecules:
Fe2+(aq) + 2H2O(l) <=> Fe(OH)2 (s) + 2H+(aq)
You can see from this equilibrium that addition of acid will cause it to adjust to the left and tend to re-instate the Fe2+(aq) ions.
Secondly, Fe2+(aq) ions are rapidly oxidised (by the air) to Fe3+(aq) ions under alkaline conditions, so the use of sulphuric acid ensures that this process is curtailed and that your solution contains the original number of Fe2+(aq) ions prior to your titration with potassium manganate(VII).
Incidentally, I would suspect that boiling your vegetables under acid conditions also causes the insoluble matter of the vegetables to break down more rapidly and release its Fe2+ content in a shorter time than would be possible using water alone.
It does not surprise me that the calculated values are “above the daily requirement”. Assuming that your calculations are correct (!), our daily requirement of mineral ions and vitamins is often met by the intake of surprisingly small quantities of vegetables and fruit. You do not mention the vegetables you have been using, but it is possible that they are ones which contain relatively large numbers of Fe2+ ions.
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updated: 24 February 2004
