Redox TitrationsWatch this thread
I've been given a redox titration worksheet from chemsheets
(I don't know if that link will work)
I've done the first 2 questions, but am confused with regards to details on question 3 and 4.
I have the mark scheme available to me because I subscribed and I can't even figure out how to do it from their working out shown.
Specifically Question 3
Calculate x in the formula FeSO4.xH2O from the following data: 12.18 g of iron (II) sulphate crystals were made up to 500 cm3 acidified with sulphuric acid. 25.0 cm3 of this solution required 43.85 cm3 of 0.01 mol dm-3 KMnO4 for complete oxidation.
I get some of it but come undone after the initial calculation of moles... Would anyone mind explaining it step by step for me?
I just cannot get my head around it
Also, Question 4....
A tablet weighing 0.940 g was dissolved in dilute sulphuric acid made up to 250 cm3 with water. 25.0 cm3 of this solution was titrated with 0.00160 M K2Cr2O7 requiring 32.5 cm3 of the K2Cr2O7. Calculate the percentage by mass of Fe2+ in the tablet.
I've done the initial steps:
- Calculated moles of K2Cr2O7
- Used the molar ratio to find the number of Fe2+ moles which i thought should have been 0.0003072
And then you'd proceed to link that with Mass = moles x Mr and work out percentage from there (comes out at approx 1.82%)
However, the mark scheme multiplies the moles of Fe2+ by 10(?); causing the final answer to be 18.25%
I don't get why? Is it to do with the volumes? E.g 250cm to 25cm?
Even still, I don't understand why....
Sorry. I realise this post sounds as confused as I am :P
Any help would be massively appreciated!
The multiplies by 10 bit is caused by the stock solution being 250 cm3 and the titration volume being 25 cm3, i.e. the stock solution contains 10x the number of mol.