Iron content in cereal

[chemrox]

I have to make a titration in order to determine the content of iron in cereal. The iron is present as iron (II) soluble salt and i have to get this out of the cereal. I was thinking about crunching the cereal, adding water and then filtering it using filter paper. Is there a better way to do this?

After this, i have to titrate it with potassium dichromate. I'm not sure whether or not we have to prepare the dichromate. That would just involve dissolving the mass into water wouldn't it? Then excess sulphuric acid would have to be added to the dichromate to obtain the Hydrogen atoms in the redox reaction.
I am not entirely sure if it has to be exces... I am worried if it will react with the iron but seeing as the iron is an ion, it shouldn't react with H₂SO₄ should it? And the SO₄ is just a spectator ion so it shouldn't interfere with the redox reaction. So the excess sulphuric acid just sits there? what does it do?

Basically, i get to make up for the figures, as long as i reach an endpoint of 25mL. The concentratin of the iron is 10mg per 60g serve of cereal... this means large samples of cereal will be needed. This is becuase the concentration of the dicrhomate shouldn't be too small as it would increase error.

[Zedekiah]

Yep, water and then filtration seems reasonable to me. I'd be wary of other substances that are going to be reacting with the dichromate. That probably won't be too big of a deal, just keep in mind that could be a source of error.

If you have to prepare the dichromate, you're going to have to measure out a precise amount on an accurate scale. Make sure you then make up the solution properly using volumetric glassware when applicable so that you know its molarity. You can acidify the dichromate solution ahead of time or acidify the unknown iron(II) solution. I've never used dichromate, just permanganate. I always acidified the permanganate when making up the solution in a volumetric flask, but I doubt it really matters. That can be in excess.

The Fe(II) isn't going to react with the sulfuric acid. You're right, it is going to spectate. See the reaction below :

Reduction half reaction :
14H(+) + 6e(-) + Cr₂O₇(2-) --> 2Cr(3+) + 7 H₂O

Oxidation half reaction :
6[Fe(2+) --> Fe(3+) + e(-)]

Overall reaction :
14 H(+) + Cr₂O₇(2-) + 6 Fe(2+) --> 2 Cr(3+) + 7 H₂O + 6 Fe(3+)

(The H(+) is coming from the sulfuric acid.)

[chemrox]

thank you heaps!!!!

that's great... now about the amount of dichromate to use... i've heard that if you have the concentration too small, the error increases...would 0.01 M of potassium dichromate be too small?

It's just, that if it needs an end point of 25mL, a large amount of cereal is needed due to the small iron content in the cereal. This poses a problem in mushing up the cereal (i think i've got something like 13kg of cereal!!!!) in order to extract all the iron.

Also, would atomic absorption spectroscopy be a better method to find the amount of iron in it or would HPLC (high pressure liquid chromatography). I initially thought AAS because it would be quicker but then realised, you would still have to have it in soluble form... still, it should be quicker than titration because all you would have to do is run it in a machine...same with HPLC too. I'm just trying to find the best method with the best reasoning.

[Zedekiah]

It is my understanding that the error increases because the weighed mass of dichromate is going to be losing sig-figs as you lose a decimal place. (IE : If I have a miligram scale and measure out 0.001 it has one sig-fig but if I measure out 0.192 it has three sig-figs.)

So it depends on the volume you're going to be making of dichromate. To be honest, I can't imagine 0.01 M being too small, but I don't know what kind of accuracy you're looking for.

I don't know much about the other methods you have suggested unfortunately.