MMendoza14
Harmless
Posts: 2
Registered: 26-6-2012
Location: State College, PA
Member Is Offline
Mood: Excited!
|
|
Iron (III) Sulfate (ferric sulfate) synthesis
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
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
'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]
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
kristofvagyok
International Hazard
Posts: 659
Registered: 6-4-2012
Location: Europe
Member Is Offline
Mood: No Mood
|
|
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]
I have a blog where I post my pictures from my work: http://labphoto.tumblr.com/
-Pictures from chemistry, check it out(:
"You can’t become a chemist and expect to live forever." |
|
|
Poppy
National Hazard
Posts: 294
Registered: 3-11-2011
Member Is Offline
Mood: † chemical zombie
|
|
Thats explains a lot why! Iron III is such a very insoluble thing to work with!
|
|
|
Waffles SS
Fighter
Posts: 992
Registered: 7-12-2009
Member Is Offline
|
|
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
Super Administrator
Posts: 7758
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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
Thought-provoking Teacher
Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline
Mood: No Mood
|
|
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
Harmless
Posts: 2
Registered: 26-6-2012
Location: State College, PA
Member Is Offline
Mood: Excited!
|
|
| 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.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
