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Determination of iron by thiocyanate colorimetry

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taa.10:
I've recently did a project where I tried to determine the iron content in a tablet by doing colorimetry and titrations. I used iron nitrate dissolved 0.505g by adding 100ml of 2M nitric acid using a 100ml graduated flask. I took 20ml of this solution and added it to a 200ml graduated flask to make a 0.001M stock solution. I prepared Fe (III) stock solutions where they went 0,2,4...12x10-5 mol/l solutions in 500ml standard flasks,With this a calibration curve was made to determine the concentration of Fe (III) from this the concentration obtained was 4x10-5. the colour change happened when I added potassium thiocyanate.
For the second part is when I actually used the tablets (each weighing 0.37g) and I took 5 iron tablets crushed them and dissolved with 100ml of the same nitric acid. Ps the tablets I used were in the form ferrous fumarate.
 I then transferred this to a 250ml graduated flask and took 25 ml of this and transferred it to a conical flask for the titrations. This was titrated with potassium permanganate. I then took 5ml of iron solution and added it to a 100ml graduated flask for samples to use in colorimetry. Starting at 2x10-5mol/l. 10 ml of 1M potassium thiocyanate was added to the solution with 2 mins in-between for absorbance readings....

This is the part where I'm confused as for the calculation part I'm not sure what equation I would use to even calculate the mass of iron in the tablet as I used so many things throughout. The experiment itself worked but I simply cannot find another experiment like this online with calculations. The only thing i can find is to use the calibration curve to get the concentration of Fe3+ to find the mass of the iron in the tablet, however for the calibration curve i did not even use tablets as this was for step 2.

I'd be so grateful if anyone would help with this as none of my teachers have even taught us how to do any calculations similar to the experiment I've done. TYSM 😭😭😭

Hunter2:
I dont understand the step with the Permanganate.

First you did a calibration curve using iron-(III)- nitrate. So far so good.
Normaly you do the same thing with the tablets. Or do the tablet contain iron-II- fumerate. permanganate is used for oxidising?
Calculation is also Lambert Beer law.
You got the calibration curve. Ideal E = εcd but sometimes an blank.
E = εcd + x
Constant are ε,d and x
So you have E~c

c has to converted to mass of iron. Then you get directly the mass of iron of your tablet. Of course consider all your dilutions.

The preparation gives 0,002 M stock not 0,001 M

0,505 g in 100 ml = 5,05 g/l.
20 ml of this solution contain 1/5 = 0,101 g
This in 200 ml = 0,101g/0,2 l = 0,505 g/l
Molar mass of water free iron-III-nitrate is 241,86 g/mol
 n= 0,505 g/241,86 g/mol = 0,002 mol/l
Or did you use the nona hydrate (403,99 g/mol) then its 0,0012 mol/l

The dissolving and dilution of the tablet get 5 * 3,7 g in 100 ml = 18,5 g/l. This solution diluted  to 250 ml = 7,4 g/l. Taken 25 ml gives 185 mg before permanganate titration, from here I cannot follow.

taa.10:
Yes you're right I should've specified that the permanganate was used to oxidise as the tablets were in the form of Fe(II). I started off by doing a relationship between the iron and thiocyanate which gave the equation Fe3++SCN---→FeSCN3.
Then as Fe3+ is turned into Fe2+ I could further use that with the redox equation between permanganate and iron equation as the ratio is 5:5 of Fe2+ and Fe3+ anyways and figure the mass from there. Would that work too?

Hunter2:
That is no need.
Fe2+ ist transferred into Fe3+
The molar mass of both are the same.
You have only to compare the absorbance from the calibration curve with your result from the oxidised iron in the tablets.

taa.10:
Okay thank you for the help.

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