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Author: Subject: Chromium isolation
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[*] posted on 27-4-2005 at 08:11
Chromium isolation


Seewwwwwww.... has anyone been tempted to make some elemental chromium? Say for the hell of it? Not much other use of it afterall.. :D

Any thoughts on purifying:
Dichromate finishes
Stainless steel
???

Tim
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[*] posted on 27-4-2005 at 19:24


How about using CrO3 from a pottery store... I know you can make an electroplating soln. with CrO3 and H2SO4. I forget the exact proportions, but it does take about 1 amp per square inch!!! :o

Electroplating of course isn't the brightest idea to isolate large amounts of anything, but that is one of the main uses of chromium... ;)

edit- oops, you are right! Cr2O3, not CrO3! Sorry.

[Edited on 29-4-2005 by Cyrus]




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[*] posted on 27-4-2005 at 21:05


You'll get Cr2O3 from a pottery place, not CrO3, and it does not easily dissolve in sulfuric or any otther acid. Fusing it with nitrates (optionally combined with carbonates) produces dichromates and chromates respectively. It may be possible to plate out metal from dichromate solution. Or you could try aluminothermic reduction of Cr2O3. That's likely to be more spectacular and less pure.



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[*] posted on 27-4-2005 at 23:17


I'd rather not try thermite, I don't have a good way to melt a relatively reactive high melting point glommed with slag.

Electrowinning is probably the most practical. IIRC, 50% of chromium production never sees a remotely metallic form, it goes right to chromic acid for plating.

I don't have a pottery store within 50 miles that has more than just clay, so that's out. :(

Tim
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[*] posted on 28-4-2005 at 00:46


The book <A HREF="http://www.sciencemadness.org/library/books/the_manufacture_of_chemicals_by_electrolysis.pdf">The Manufacture of Chemicals by Electrolysis</A> indicates that dichromates and permanganates can be made by using chromium or manganese-bearing iron alloys as the anode in electrochemical cells with alkali hydroxide solution as the electrolyte.

Quote:
E. Lorenz (2) has shown that it is possible to produce permanganate by electrolysing a solution of caustic potash, if a manganese or ferro-manganese anode be used and a cathode of copper oxide (the positive plate of a cupron cell for example). The same method can be used for preparing potassium bichromate if the anode be of ferro-chromium. In both cases the iron in the anode is converted to ferric hydroxide which collects at the bottom of the cell.

(2) Zeitsch. anorg. Chem., 1896,12, 393, 396.


You might try using a stainless steel as the anode in a cell with NaOH solution and see what you get. The iron may provide plenty of gunk to complicate things. See what goes into solution.

Edit: linkified book and quoted relevant passage.

[Edited on 5-1-2005 by Polverone]




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[*] posted on 29-4-2005 at 04:41


You could isolate it by what this guy calls "reverse electroplating"

http://yarchive.net/car/reverse_electroplating.html

I looked for this because I remember an uncle doing something like it when I was a kid. It cleaned his old Chevy car part in any case. Oh how that man loved his car.




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[*] posted on 30-4-2005 at 03:30


" IIRC, 50% of chromium production never sees a remotely metallic form, it goes right to chromic acid for plating. "

and then....?

If you use a slight excess of Al when reducing Cr2O3 then the slag should dissolve in NaOH soln leaving the Cr behind.
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[*] posted on 30-4-2005 at 04:11


Remember that Chromium is also amphotheric, like aluminium, and will react with alkalis (although not with the same vigour IIRC). A thermit reaction will be difficult to control. I once carried one out. I must admit it was spectacular, but the whole thing spread out the chromium which evolved around the spot, barely anything remained in the crucible.

[Edited on 30-4-2005 by Esplosivo]




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[*] posted on 30-4-2005 at 04:20


Like I said, I'd rather not do thermite. Plus I have no good way of melting the result into a homogenous blob.

Electrowinning seems to be the way to go, just need to get it in solution. I'll try zapping some stainless today.

Tim
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[*] posted on 30-4-2005 at 04:58


Well, I got once some chromium globules that were definately made via a termite reaction. There was some yellow stuff left on the globules aswell.
Using less Al then necesary is the key to a slower reaction.
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[*] posted on 30-4-2005 at 07:39


I've been electrolzing a strong lye solution between two stainless electrodes (at least that ought to be what they are, too heavy for aluminum, which would be long gone in the electrolyte anyway, and nonmagnetic so aren't plain steel), last check a thin gray deposit was growing on the cathode but the solution is only slightly yellow (probably more from soaking some dichromate finished hardware with the solution first, which only blackened them). So far it seems to be a good Brown's gas production cell. :(

Edit: the gray layer is up to about 1/8". The anode doesn't appear to be plating out, but its surface is discolored orange/red. The solution is still only very slightly yellow.

Tim

[Edited on 30-4-2005 by 12AX7]
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[*] posted on 30-4-2005 at 13:34


"Remember that Chromium is also amphotheric, like aluminium, and will react with alkalis (although not with the same vigour IIRC)"
That's news to me and to those who wrote the text books I have seen.
Granted, if you were to oxidise it to Cr(VI) it would be, but that's not really relevant to washing the metal with caustic.
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[*] posted on 30-4-2005 at 13:57


To my surprise, when I tried it with a stainless steel spoon as anode, and a steel nail as cathode, I immediately began to see purple/red swirls coming from the anode. This happened when I switched the polarity too, mostly near the bottom of the cell where some undissolved NaOH remained. It seems that I must have been producing ferrate? I don't know of anything else with that color that could come from a steel nail.

After a couple hours of running the cell (5v, 5 A max (it got pretty hot)), I started pipetting solution out and introducing it to various other compounds. Ascorbic acid, ethanol, salicylic acid, sucrose, and hydrogen peroxide all seemed to reduce (?) it to an unremarkably iron-yellow color. Acids seemed to discharge the color rapidly as well.

I also tried seeing what happened when I ran the cell with an oxalic acid solution, the same spoon at the anode, and the nail at the cathode. I obtained a solution of very interesting color. In small amounts, like in a pipet, it appeared pale green like a solution of iron (II) compounds. In larger amounts or different lighting, it looked red. It seemed there was some very fine dispersion of particles that made the solution appear red in certain lighting conditions while the "true" transmitted color was green.

None of this puts you closer to getting pure chromium from stainless steel, but I still found it interesting.




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[*] posted on 30-4-2005 at 17:09


Hmm odd.

Well the cathodic growth is up to 3/16" thick, I'll scrape it off and see if it...uh, well no response to bases obviously, but I can try heating, acid, etc...

I'll try acid and neutral (salt) later.

Tim
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[*] posted on 30-4-2005 at 17:33


Cr<sup>3+</sup> + 4OH<sup>-</sup> -> Cr(OH)<sub>4</sub><sup>-</sup>

This reaction drives the equilibrium between solid Cr<sub>2</sub>O<sub>3</sub> and the dissolved form to the left, dissolving the solid.
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[*] posted on 1-5-2005 at 06:23


I just wonder if you read my post?
What did you think was the point of adding excess Al? Did it occur to you that there shouldn't be much Cr2O3 left over?

It's generally the case that oxides get more acidic as the oxidation state rises. Cr is a case in point; CrO3 is strongly acid, Cr2O3 rather weakly so CrO isn't and (and this is the important bit) unlike aluminium, chromium as the metal (which is what this thread's about) isn't soluble in alkali.
So, if you did this thermite reaction (and I can understand someone not wishing to, it's quite violent) you would end up with Chromium (goody goody!), Al2O3, Some Al, some Cr2O3 and a few odds and ends from the crucible.
Al2O3 disolves in alkali because it's amphoteric
Cr2O3 disolves in alkali because it's amphoteric

Al disolves in alkali because its one of the amphoteric metals (there aren't many)
And the Cr sits there because it doesn't disolve in alkali. It would do if you left it long enough in the presence of oxygen but, in this (ie the relevent) case its not going to because of the Al disolving and scrubbing any O2 from solution.

Chromium sesquioxide is amphoteric. You could just about argue that the trioxide is too but you would be pushing it to find an instance of it acting as a base. Chromium, a metal, isn't amphoteric.
BTW, you haven't said what drives that eqm to the right.

[Edited on 1-5-2005 by unionised]
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[*] posted on 1-5-2005 at 06:36


Sorry, I thought the post was referring to dissolving the oxide, not the metal. I agree, the metal shouldn’t dissolve in a base with any significant speed.
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[*] posted on 5-5-2005 at 11:52


BUMP!

Am currently electrolyzing stainless (food grade I think, 304 or 316?) in HCl (half dilute muriatic acid) at 5V. Strong green after a minute or two.

Should go check it in case something's plating out and growing towards a shorted cell...

Edit: That went smoothly. I now have a (relatively neutralized with NaOH, no H2 bubbles from Zn metal testing) solution of Fe + Ni + Cr + Cl. Now, to seperate!

Something interesting I noticed was the range of boiling/sublimation points, CrCl3 (- is CrCl2 formed, or is that less preferred?) seems to be the highest. I could do worse than loading the mixed, dried salts in a copper tube, heating just shy of melting the Cu and scraping out the sublimation layers. ;)

Anything else? Precipitate (say with Na2CO3) then basify with lye solution, bubbling air to oxidize to chromate? Then finally of course plating out a nice shiny gob of the stuff. :)

Tim

[Edited on 6-5-2005 by 12AX7]
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[*] posted on 6-5-2005 at 21:23


Say, instead of atmospheric oxidation, would bleach (Ca(OCl)2 to be exact), or sodium chlorate for that matter, oxidize it at all? A small test (without basification) certainly "bleaches" the iron to ferric hydroxide quite nicely. :P

Oh, was going to add, I precipitated it all with sodium hydroxide and carbonate. So there's a mixture of hydroxides and carbonate precipitates (a shaly, greenish pastel blue), well whatever...

Tim

[Edited on 7-5-2005 by 12AX7]
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[*] posted on 6-5-2005 at 21:48


Quote:
Originally posted by unionised
"Remember that Chromium is also amphotheric, like aluminium, and will react with alkalis (although not with the same vigour IIRC)"
That's news to me and to those who wrote the text books I have seen.
Granted, if you were to oxidise it to Cr(VI) it would be, but that's not really relevant to washing the metal with caustic.


Isn't chromium metal amphoteric? I mean not to the extent of aluminium, but if boiled with a concentrated NaOH solution shouldn't it react? If not then I'm sorry for confusing you, but damn it its written on my class notes. I hate these teachers.




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[*] posted on 7-5-2005 at 02:18


I haven't tried it (though, if I find some chromium about the place I might) so I can't be sure.
Even if the oxidation potential favours it, which I doubt, in order to disolve it would have to get oxidised to Cr(III) and I think it wouldn't get past the insoluble Cr(II) without an oxidant, like air, present.
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[*] posted on 9-5-2005 at 12:28


Added lye to the "stainless hydroxide" mixture (which BTW has been producing floating brown material, I'm guessing iron (in preference to Cr(III) and Ni(II)) is oxidizing to ferric oxy/hydroxides.

Anyways, I'll tell you what color the solution is once the Ni & Fe ppt's settle.

In the mean time, can I reduce chromates with sugar, possibly slightly acidified? I don't have any ferrous sulfate on hand, and haven't been able to make any (I made a 50% sulfuric acid (liquid fire) solution and added rust and metal with no reaction!?). Or can I just flush it without guilt?

Tim
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[*] posted on 11-5-2005 at 15:48


Solution has remained perfectly colorless after several days of soaking with lye and air exposure (lye's probably weakening from CO2 absorption now! Say, is CrO4 stronger than CO3?).

I am officially out of ideas. Anyone?

Tim
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[*] posted on 11-5-2005 at 17:52


The only way I know of to separate the metals at this point would be to add NaOH/peroxide to get soluble chromate or filter and roast in air with or without oxidizer with carbonate and/or hydroxide and leach the chromate.
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[*] posted on 19-5-2005 at 03:49


I have dried the iron, chromium and nickel hydroxide/carbonate precipitates and calcined to a red heat, about 1400-1500°F (750-800°C). (Geez, now I'm listing temperature in *three* units?!)

I put it back in a lye solution and at last, it is turning yellow. :D Any ballpark as to how long I have to wait for it to dissolve? Would hypochlorite speed it up?

Tim
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