12AX7
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Copper-Iron Mixtures
I've been processing this solution of copper chloride with various impurities. I've got most of the stuff out, but I can't manage to crystallize the
last part. When I dehydrate it, it forms a green mush, as if it were copper chloride, but something is keeping it from forming good crystals; or
else, something else is co-crystallizing, as if it were an eutectic.
The weird thing is, after handling the solution, on rinsing my glassware I get an orange/yellow solution, or possibly a colloidial suspension. This
is after like 1:20 dilution, so for the color to be as strong as it is, the impurity must be very darkly colored.
Finally, this evening I added some soda to the stuff. I discovered that the solution is remarkably acidic, much more so than residual HCl would
allow. A "weaker" solution, obtained by dissolving a pale residue left from dissolving this indeterminate crystallized mass, finally turned brown
after much soda was added. The curious part is that, acidic, the solution was light green, while more basic, it turned deep red/brown, without any
obvious precipitate. After more base it finally gelled into a thick precipitate, suggesting quite a lot of Fe(OH)3.
Might I have a copper chloroferrate(III) sort of combination here, or what?
Tim
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not_important
International Hazard
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Iron seems to be a common impurity in copper compunds. In the case of the sulfate Fe(II) is isomorphic with Cu(II) and it is next to impossible to
separate them through crystallisation. I don't know about the case of the 2+ chlorides.
If you check some of the older inorganic prep books you'll often find one or two preparations of pure copper salts. I think the trick was to
have a slightly acid solution, add an oxidiser such as HNO3 or H2O2, heat for awhile, then adject the pH with ammonia or fresh Cu(OH)2 (in excess).
Cool, filter to remove ferric hydroxide. Check for remaining iron by taking a small sample of the solution, boiling with a few drops of HNO3 or a bit
of Na2O2, adding strong aqueous ammonia, spotting it onto filter paper, and washing the spot with more ammonia solution. The copper is washed away,
iron will leave a yellow-orange-brown spot. If that happens you repeat the oxidation and pH adjustment of the main solution. Check in those inorganic
prep books, it's been a long time and I'm not sure I'm remembering correctly.
You're never goingto be able to get really pure materials with 100% recovery. A lot of the impurities, both anion and cation, are going to be in the
last few percent of mother liquor. It's better to not try to pull everything out of it, but just evaporate it separately and toss the solids back
into the starting point of another batch of whatever it is you're making (for ordinary salts, not tricky things). Even then, if the impurities are
something that is not removed in the normal process steps, you'll reach a point where you have to discard that last 5 or 10 percent of the batch; in
some cases it's worth working it up for the impurities - potassium salts sometimes have some Rb and Cs in them for example.
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unionised
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IIRC you can extract FeCl3 from solutions containing excess HCl by washing with ether. Might help. I can't see you separating Fe(II) from Cu(II)
nearly as easily as from Fe(III). As not important says Fe(II) and Cu(II) tend to be isomrphous.
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12AX7
Post Harlot
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I'm familiar with Fe(II) being with, lemme see here, Mg, Zn, Mn(II), Co(II), Ni(II), and I think a few others. Any can replace magnesium in spinel
(MgAl2O4) and many other compounds. But I'm not familiar with a Cu(II) spinel crystal, though it's quite close to Ni(II) which is quite close to
Fe(II).
At any rate, the solution was acidic and exposed to air for quite some time -- enough to make it goodly Fe(III), I would think. This isn't supported
by the color, but it IS supported by the response to base. Go figure!
At the moment I have one solution, the "weak" one, filtering some drab orangey sort of precipitate -- most likely Fe(OH)3, probably with Cu(OH,Cl)2 or
something. There are still globs of browner FeOOH or whatever it is. The other solution, which is not filtering at the moment, has a greener color,
likely due to Cu(OH,Cl)2 from adding too much soda. I'll filter that and wash it with HCl, or maybe I'll get a jug of 10% NH3. I need a trip to that
aisle anyway...running low on HCl and H2SO4.
Tim
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not_important
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The problem isn't with the formation of a compound or double salt, it's just that the ion sizes are similar to each other that one can randomly
replace the other.
I don't think that simple exposure to air is enough to make sure all the iron is in the F3(III) state. They always seemed to add the oxidiser, HNO3 or
H2O2, and then bring it to a boil for a few minutes. Then the base would be added while keeping the solution hot, to encourage the formation of a more
filterable precipitate. They often allowed the ppt. just settle out, then tested the supernatant liquid for complete precipitation of iron. Finally
the supernatant solution would be decanted through a filter.
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