bbartlog
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Cu and Fe sulphates... what happened here?
Starting with a solution of MgSO4 (approximately one molar), I did some electrolysis using a flowerpot as an ion exchange membrane and a copper pipe
cathode, with a view towards producing some Cu2SO4.
The flowerpot tends to leach iron (and possibly small amounts of other elements), so after a couple of days I ended up with a blue-green solution with
I believe to have been a mixture of Cu(II)SO4, Fe(II)SO4, and MgSO4. (Also a flowerpot full of Mg(OH2), which I had no use for).
I began boiling this solution to concentrate it (wanted to see what would crystallize). It
- quickly darkened
- small flecks of what looked like elemental copper were deposited at the waterline
- small amounts of flocculent orange-brown precipitate formed
- the solution became less greenish and more blue (once the floc had settled)
My assumption is that I oxidized Fe(II) to Fe(III), in the process reducing the copper. As I understand it, cuprous sulfate Cu2SO4 is unstable and
will decompose to Cu + CuSO4. This would also explain the color changes in the solution. However, what is the other precipitate? Fe2(SO4)3 is supposed
to be soluble. The MgSO4 should be stable here, I think it plays no part in the reaction (but maybe it affects the solubility of other compounds?).
I see another reference that suggests an insoluble iron compound can be produced in a similar situation. The encyclopedia of chemistry (from 1862)
says:
'As a mixture of copper and iron vitriol is often employed in the arts, it is not always necessary to obtain blue vitriol free from iron; [...]
the sulphates heated to beginning redness in an iron vessel or reverberatory, will leave nearly all the iron salt insoluble, together with a
little of the copper'
Aside from the reaction conditions, sounds about like what I saw - but they don't describe the reaction products either :-(.
Is this some mixed compound, e.g. Cu(I)Fe(III)(SO4)2? What am I seeing?
I still have a couple grams of the orange-brown precipitate, if anyone has suggestions on tests that could be done on it. As a first step I'm going to
figure out whether it's actually a single compound and not just a finely divided mix of say Cu2O and Fe-who-knows-what...
[Edited on 9-9-2009 by bbartlog]
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12AX7
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Cu(I) would fit the description, but it's rather difficult to produce in solution. Typically, electrolysis of a copper anode (with no membrane) in a
solution where Cu(I) is fairly unstable (nitrate, sulfate, etc.) leads to Cu(OH)2 precipitate (or CuO at higher temperatures). The anode prefers to
oxidize directly to Cu(II), possibly with an interfacial layer of Cu2O.
Only in complexing environments, such as a concentrated chloride solution, is Cu(I) stable enough to precipitate as yellow to brick-orange Cu2O on
reaching the cathode's alkalinity.
Mind that Fe2O3 can look very much like Cu2O, depending on the relative particle sizes.
I'm not sure why you would think iron was dissolved. At the anode, oxidation should occur in direct proportion to the amount of acid drawn towards
it. There's no reason for iron to dissolve. If it did, you may consider soaking your pot in hydrochloric acid (change the acid until it no longer
turns green) to remove the iron first. Or, you may consider a low iron material, like plaster of paris. Calcium sulfate will be stable in a sulfate
solution.
The quoted passage refers to the easy decomposition of iron sulfate, while copper sulfate is more stable. Therefore, sulfur fumes (SO2 and SO3 gas)
are released by the iron, while CuSO4 remains, which can be leached out of the product, leaving iron behind as Fe2O3. The conditions are completely
different and not applicable.
Tim
[Edited on 9-9-2009 by 12AX7]
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bbartlog
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I'm not sure why you would think iron was dissolved.
I'm assuming dissolved iron, partly because of the color of this solution (greenish tinge prior to this reaction, partly because other similar
electrolysis I've done (with e.g. Na2SO4) has resulted in a yellow-tinged solution without the low pH that a solution of H2SO4 or NaHSO4 of sufficient
concentration should have had. And Fe(II)SO4 seems to have the right color. And while the iron oxide used to color the flowerpot seems fairly
insoluble, there's an electrical field and SO3- floating around which certainly makes mobilization plausible...
Mind that Fe2O3 can look very much like Cu2O, depending on the relative particle sizes.
I'm leaning towards this as the precipitate. Fe(III)2(SO4)3 seems unstable enough that a couple hours boiling would decompose a little of it (IIRC
FeCl3 will give off chlorine gradually if boiled in solution), with the iron oxide remaining as a residue. Cuprous anything doesn't seem like it
should persist in this environment. And the color looks a lot like rust... boring after all :-).
you may consider a low iron material, like plaster of paris.
This is a great suggestion! I have a bunch of plaster of paris and could cast it as desired; hadn't really thought to use it before, figuring that it
would disintegrate, but as you say it should be stable in sulfate...
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chloric1
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not to mention that plaster of paris is porous enough to have a clay slurry poured into it so to absorb excess water and leave the solid clay behind
in whatever shape the plaster mold is in! It may be TOO porous as a cell divider
and you may need to impregnate it with sodium silicate solution prior to use.
Fellow molecular manipulator
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bilcksneatff
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I would think making copper(II) sulfate this way would be kind of hard, because it seems like removing all of the MgSO4 would be difficult. I think
CuSO4 is about twice as soluble as MgSO4 in boiling water, so you may be able to do that.
Another option (to avoid the flowerpot) would be to use some kind of salt bridge instead.
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watson.fawkes
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Quote: Originally posted by 12AX7  | | I'm not sure why you would think iron was dissolved. At the anode, oxidation should occur in direct proportion to the amount of acid drawn towards
it. There's no reason for iron to dissolve. |
A membrane (the flowerpot, presumably terra cotta) with an an
electrolytic current across it creates something like a virtual anode on the cathode side (and vice versa on the other side). Assuming we have protons
as our main charge carrier, that means our virtual anode, not considering materials, is streaming out protons. It's completely believable that some of
this current is reducing iron at the surface, at some rate, so that the virtual anode is streaming out a mixture of hydrogen and iron ions.
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12AX7
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That's true, it's possible that a mixture of ions is coming from the pot. Doesn't really matter as long as charge is conserved, eh?
Tim
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