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I am carrying out a colorimetric analysis into the amount of iron in spinach after it has been cooked for varying amounts of time. I need to make an iron(III) solution in order to produce my calibration curve. It has been suggested that I use Fe(NO3)3.9H2O but will need to standardise it as it is not a primary standard. How do I standardise an iron(III) solution or is there a better Fe(III) compound to use for my standard?
Igloo writes ...
First see if you can find out whether the double iron(III) salt, (NH4)2SO4.Fe2 (SO4)3.24H2O is a primary standard. Double salts often are. If so, go for this one – it’ll make your life a lot easier! If it isn’t, or if it is unavailable to you, then, yes, I would also choose Fe(NO3)3.9H2O, though you could use FeCl3.6H2O instead. In either case the solution needs to be made up in acidic conditions in order to minimise hydrolysis to Fe(OH) 3.
Although you could standardise this solution with a standard solution of potassium dichromate(VI), via reduction with zinc, dichromate(VI) titrations are hard to carry out accurately, so I would do the following, cumbersome though it may be.
Start by making up a standard solution of ammonium iron(II) sulphate, (NH4)2SO4.FeSO4.6H2O. This
compound, also known as Mohr’s salt, is certainly a primary standard and should be available to you. Also make a solution of potassium manganate(VII) of appropriate concentration and carry out titrations with acidified portions of the iron(II) solution to determine the exact concentration of the manganate(VII) solution. Secondly pipette portions of your iron(III) solution into a flask, add an excess of acid (e.g. dilute H2SO4) and some granulated zinc, and wait until the reduction to Fe(II) ions is complete; then titrate with your manganate(VII) solution. This will enable you to determine the accurate concentration of the Fe(III) ions in the iron(III) solution.
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Risk assessment
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updated: 17 December 2006
