Sciencemadness Discussion Board

Iron (III) Sulfate (ferric sulfate) synthesis

MMendoza14 - 26-6-2012 at 12:58

Hey everyone! I'm having troubles producing iron sulfate... I must start with ferric oxide and use sulfuric acid to produce iron sulfate but ive been having some problems... Ive tried using 0.1 M and 1 M sulfuric acid and I'm thinking about making the jump up to 5 M in hope of getting better results. So far I've produced some tannish colored solid on top of unused hematite, even when I'm using proper stoichiometry given the sulfuric acid concentrations. I've tried ranges of starting material from 0.5 g all the way to 2 grams without much success... After I react the two together I put the solutions into the oven at 120 C for however long it takes for the liquid water to be removed. Whenever I test the pH of the supernatant I find that its always very low, implying to me that the sulfuric acid has not reacted very much. There's also the problem of ferric sulfate being soluble in water which makes removal of the acid kind of difficult... I hope one of you guys can help!! Thank you!!

Hexavalent - 26-6-2012 at 13:18

'I must start with ferric oxide' - is this a school project, or an amateur thing? How flexible are the rules?

Do you stir well when you add? How pure is the iron (III) oxide? What form is it in - granules, powder etc.?

Have you tried adding it to hot sulfuric acid?

[Edited on 26-6-2012 by Hexavalent]

kristofvagyok - 26-6-2012 at 13:45

Use a bit more concentrated sulfuric acid, 1M acid is not even good for washing my hands if it get's dirty in the lab....

5-8M acid will do it, esperially if you heat it. Or just add some H2O2, it will do the rest(:

[Edited on 26-6-2012 by kristofvagyok]

Poppy - 26-6-2012 at 17:30

Thats explains a lot why! Iron III is such a very insoluble thing to work with!

Waffles SS - 27-6-2012 at 00:19

I have experience about Iron Components, Iron(III)Oxide is really stable and Nitric acid and Sulfuric cant dissolve it(just Hot Hcl can do it)
(see below topic about Ferric Nitrate)
http://www.sciencemadness.org/talk/viewthread.php?tid=17161#...

You can easily react Iron wool or Powder with warm %50 Sulfuric acid (this make Iron(II)Sulfate) and then add Hydrogen Peroxide(or nitric acid or even chlorine gas) for oxidation Fe2+ to Fe3+

Fe + H2SO4 → FeSO4 + H2

6 FeSO4 + 3 H2SO4 + 2 HNO3 → 3 Fe2(SO4)3 + 4 H2O + 2 NO(beware of mother fu...ker NO and NO2 gas)
or
6 FeSO4 + 3 Cl2 → 2 Fe2(SO4)3 + 2 FeCl3
or
12 FeSO4 + 3 O2 → 4 Fe2(SO4)3 + 2 Fe2O3

I advise first reaction because by Hydrogen peroxide or air oxygen as oxidation agent you will get Iron(III)Oxide too and this is hard to get rid of it.



[Edited on 27-6-2012 by Waffles SS]

woelen - 27-6-2012 at 01:12

Commercially available red/brown ferric oxide (Fe2O3) is amazingly inert and the same is true for the black oxide Fe3O4. I have both oxides from a pottery supplier and they do not dissolve appreciably, not even in hot 50% H2SO4 or boiling HNO3. It does dissolve in hot concentrated HCl, but only very slowly and a very large excess amount of acid is needed to get all of it dissolved. The solution in HCl becomes bright yellow, due to formation of the FeCl4(-) complex.

blogfast25 - 27-6-2012 at 04:25

Quote: Originally posted by woelen  
Commercially available red/brown ferric oxide (Fe2O3) is amazingly inert and the same is true for the black oxide Fe3O4. I have both oxides from a pottery supplier and they do not dissolve appreciably, not even in hot 50% H2SO4 or boiling HNO3. It does dissolve in hot concentrated HCl, but only very slowly and a very large excess amount of acid is needed to get all of it dissolved. The solution in HCl becomes bright yellow, due to formation of the FeCl4(-) complex.


Yep. My experience too. Try fusing with Na or K bisulphate...

MMendoza14 - 27-6-2012 at 08:34

Quote: Originally posted by Hexavalent  
'I must start with ferric oxide' - is this a school project, or an amateur thing? How flexible are the rules?

Do you stir well when you add? How pure is the iron (III) oxide? What form is it in - granules, powder etc.?

Have you tried adding it to hot sulfuric acid?

[Edited on 26-6-2012 by Hexavalent]


I suppose I should have been more specific with my intentions... I'm attempting to model the process of sulfur species as a pollutant and then ultimately sulfuric acid and their effects on mineral dust in the atmosphere. Ferric Oxide (hematite) is a common mineral dust particle. I am using 99% pure hematite powder. Due to the attempt to model actual atmospheric conditions I'm already stretching a little bit using 1M as the pH never really reaches that low with a minimum usually around 3-4. Many of my experiments have shown that hematite does seem to be inert, but a reaction seems to occur after I put the solutions in the oven to dissolve the excess water. Before the oven there's the same red color and clear acid and after the oven theres a brownish solid leftover implying heat may cause the reaction to proceed but the product is unidentifable using ATR and XRD so far.