ludemas19
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Oxidation of iron to iron(II/III) with copper sulfate
Good evening, for another project I need some iron (III) sulfate. I'm starting off with some common iron nails and, as oxidiser, some copper sulfate.
Apparently, the reaction proceeds to some degree and some amounts of iron (III) could be observed (yellowing of the solution, some flakes that looked
like iron(III) oxide...). However, I have no way to be sure of how much of the iron in solution is in the 3+ state rather than the 2+.
I'll briefly explain the procedure:
-Added in a test tube a few milliliters of commercial HCl solution (10-ish%)
-Added copper sulfate (a greenish coloring was observed, due to the formation of Cu-Cl complexes)
-Added iron nail pieces
-Very gentle heating (brought some water on the verge of boiling and then let the test tube rest in there)
At first discoloring of the copper solution was observed, accompanied by formation of copper metal precipitate. After the heating the color got
slightly yellow. After a few hours, and the solution cooling down, the color went back to colorless and finally to a pale blue.
At the end, some traces of iron (III) were still present (another experiment somehow proved it: it made a black complex with tannins, something that
iron(II) cannot do).
My guess on this is that chloride ions can mess up the equilibria a bit, as far as i know they tend to make complexes with copper(II) and iron(III).
Furthermore, iron(III) chloride is used to etch copper circuits, so I think it can easily oxidise back copper metal.
How could I optimize this method? Ideally, I'm thinking of phasing out HCl and substituting it with acetic acid. Maybe, I should add it later on, as
the reaction ends, because I just need it for the acidic condition to keep the iron(III) in solution.
As for oxidant, I'd still keep the copper sulfate, also because the procedure for the tannin complex requires that (not completely sure).
Any suggestions? Thanks in advance.
I'm attaching the picture of the tannin experiment: I soaked a paper strip in tea (rich in tannins), dried it and then spotted it with the iron
solution. The black spots should be the complex (on the left a control, unspotted paper strip).
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vertexrocketry
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dissolve 1 gram of potassium thiocyanate in 10 mls of water then add 2 mls of your uncertain iron sulfate solution
there is iron(III)sulfate if it turns a blood red colour
potassium thiocyanate is easily found on etsy
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j_sum1
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You need an oxidiser to get iron in the +3 state. Air is often sufficient. Hydrogen peroxide is going to be quicker and more reliable.
If sulfates are your target then avoid any other anions. IOW use sulfuric acid not hydrochloric.
For similar reasons, remove copper from your reaction. You do not need the copper or its complexes.
This leaves you with, iron nails, hydrogen peroxide and sulfuric acid.
If you do not have sulfuric acid then I guess you are stuck with HCl. Use epsom salts as your sulfate source. That will give you a mix of Mg2*, Cl-,
Fe3+ SO4-- in your solution. The least soluble ionic compound possible in this mix is your desired Fe2(SO4)3 which should precipitate out if the
solution is sufficiently concentrated. Yield will be lower but with such inexpensive reagents, this is hardly a problem. Just remember that Epsom
Salts is the heptahydrate, MgSO4.7H2O. Factor this into your stoichiometry
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bnull
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If you're worried about copper in solution and if all iron is +3, try this: use sodium carbonate to precipitate copper and iron hydroxides/oxides, add
enough sodium hypochlorite or hydrogen peroxide to oxidise all iron from +2 to +3, wash the precipitate a few times with water, shake the precipitate
with a solution of ammonia (ammonium hydroxide, ammonium carbonate, ammonium bicarbonate, it doesn't matter) to dissolve copper, filter and wash a few
times with water, then dissolve the iron sludge with the acid of your choice.
As far as I remember, the important part is iron (iii), the anion doesn't matter much as long as the ferric salt is soluble.
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Texium
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Thread Moved
6-2-2025 at 22:11 |
ludemas19
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Quote: Originally posted by j_sum1  | You need an oxidiser to get iron in the +3 state. Air is often sufficient. Hydrogen peroxide is going to be quicker and more reliable.
If sulfates are your target then avoid any other anions. IOW use sulfuric acid not hydrochloric.
For similar reasons, remove copper from your reaction. You do not need the copper or its complexes.
This leaves you with, iron nails, hydrogen peroxide and sulfuric acid.
If you do not have sulfuric acid then I guess you are stuck with HCl. Use epsom salts as your sulfate source. That will give you a mix of Mg2*, Cl-,
Fe3+ SO4-- in your solution. The least soluble ionic compound possible in this mix is your desired Fe2(SO4)3 which should precipitate out if the
solution is sufficiently concentrated. Yield will be lower but with such inexpensive reagents, this is hardly a problem. Just remember that Epsom
Salts is the heptahydrate, MgSO4.7H2O. Factor this into your stoichiometry |
Thanks, I'll try. Looking for the epsom salt, sadly online i'm only finding overpriced feet cleaner salts (I'll admit, I'm quite new to getting
reagents on my own, so maybe i'm doing something wrong), so I think i'll be going with plant fertilizer. I've checked it contains 30% potassium oxide
equiv. 10% magnesium oxide equiv. 42% SO3 equiv. As long as i need the sulfates i assume it should be good?
For the crystallisation, I should perform the reaction and later let it rest? Would you suggest cooling or would that risk too much contamination from
the other salts?
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chempyre235
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If you're in the US, Iron (II) sulfate, also called ferrous sulfate, is commonly available as well. It is used as a fertilizer. Like Epsom salt, it
will also be the heptahydrate. It too is very inexpensive and seems to be very water-soluble.
[Edited on 2/7/2025 by chempyre235]
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