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DraconicAcid
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I'm confused. Surely the easiest way to separate chromium from iron, nickel, and manganese is by precipitating the other metals with conc. sodium
hydroxide? Chromium forms a soluble hydroxy complex, which can be easily oxidized to chromate, while the others form insoluble hydroxides.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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JJay
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I think that could work but might require using a large excess of sodium hydroxide and strong stirring to avoid precipitating the chromium along with
the other metals.
[Edited on 23-6-2017 by JJay]
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Fulmen
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Keep talking... I have plenty of sodium hydroxide. Fusing with sodium hydroxide should speed up the process significantly.
I chose the direct chromate route as it should be simple to reduce it if needed. Starting at the top in one simple step sounded like the most
efficient approach. If it works, that is.
And the chemistry supports it, right? Ferrates are stronger oxidizers than chromates, and manganates decomposes with heat. Right?
Or are we in equilibrium territory?
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DraconicAcid
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Quote: Originally posted by JJay  | | I think that could work but might require using a large excess of sodium hydroxide and strong stirring mixing to avoid precipitating the chromium
along with the other metals. |
If the chromium hydroxide precipitates, you just add more sodium hydroxide.
This is the standard way to separate chromium from other metals for qualitative analysis. Make it basic, heat with hydrogen peroxide, centrifuge out
the Fe(OH)3, MnO2, Ni(OH)2 and other crud. You may get contamination by zinc or aluminum, but I don't know how common those are in stainless steel.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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JJay
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Seems reasonable.
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JJay
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| Quote: Originally posted by Fulmen |
And the chemistry supports it, right? Ferrates are stronger oxidizers than chromates, and manganates decomposes with heat. Right?
Or are we in equilibrium territory? |
Ferrates are stronger oxidizers than chromates... do ferrates oxidize chromium (iii) to chromium (vi)? Manganates typically decompose with heat....
I'm not sure if there is an equilibrium with chromium(vi) hydroxide or not, but I remember noticing issues dissolving chromium (iii) hydroxide in a
roughly stoichiometric amount of sodium hydroxide until peroxide was added, causing it to dissolve immediately.
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subskune
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| Quote: |
Surely the easiest way to separate chromium from iron, nickel, and manganese is by precipitating the other metals with conc. sodium hydroxide
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If that works this not cost effective chromium preparation (hcl is not the cheapest, the other reactants are even more expensive) could be turned into
a cheap chromium source from waste stainless steel. I have an old pot I used for boiling anything which is now ready for the waste container. My idea
is to fill it with conc. nacl solution and use the pot as anode while adding its handle as cathode in the center. This should dissolve the pot as
chlorides and later precipitate those as hydroxides. Once the pot is done some extra drain opener could be added to dissolve the chromium and voila
chromium extracted. Thats it so far for armchair chemistry.
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S.C. Wack
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| Quote: Originally posted by DraconicAcid |
If the chromium hydroxide precipitates, you just add more sodium hydroxide.
This is the standard way to separate chromium from other metals |
* But not substantial quantities of Fe
** After prior removal of Fe
H2O2 and base will mostly make oxygen, but a little will survive long enough to make rust and chromate. Then you'll need a good separation for Cr from
much alkali.
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Fulmen
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Huh. After drying and igniting the precipitate/filter residue the extract is coming out orange rather than yellow. I tested with a pinch of sodium
carbonate, no reaction. Wouldn't that rule out both iron salts and dichromate? pH is anybody's guess, the strips doesn't survive the excess bleach.
Could be manganates, boiling should be revealing. I can't imagine ferrates or manganates would survive that for long.
That brings up another question, chromate or dichromate as the final product? Sodium dichromate is much more soluble, which should aid in the removal
of sodium chloride. It can then be treated with sodium hydroxide to reduce the solubility. Or perhaps ammonia, creating ammonium chromate. Right? Or
are they too close to the sodium salts to provide separation?
Reducing it back to Cr(III) could open up new routes to a pure product, but I don't want to go there unless I absolutely have to.
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JJay
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I've seen that before too... the equilibrium between chromate and dichromate is pretty sensitive to small changes in pH and concentration, and I
imagine it is affected by other ions in solution. Sometimes a bright orange solution will produce a bright yellow precipitate. I don't know what's
causing your solution to turn orange, but one possibility you might consider is dissolved carbon dioxide.
I think dichromate is usually preferable to chromate, but it does depend on what you think you might be doing with it.
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Fulmen
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I see. I'll work from the assumption that it's just dichromate shining through.
I agree that dichromate is the obvious end product, but sodium dichromate is so soluble it will be hard to crystallize out cleanly. Potassium
dichromate is the obvious alternative.
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JJay
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I did find it to be fairly hard to crystallize out sodium dichromate uniformly on a 10-30 gram scale, but crystallization of sodium dichromate
dihydrate is a common industrial process that has been in use for more than a century, and I suspect it's easier with hundreds of grams of material.
On a small scale the process needs to be closely monitored, and that requirement is relaxed somewhat with a larger quantity of material. A dessicator
would be required to dry the product. For many purposes, the high solubility of sodium dichromate is desirable, but it is deliquescent and certainly
harder to obtain in pure form than potassium dichromate, and the filtrate from a sodium dichromate crystallization would surely contain a lot of
hexavalent chromium. Of course, waste hexavalent chromium solutions can be recycled and certainly have their uses....
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Fulmen
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Unfortunately I seem to have hit a dead end there myself. I've boiled down the solution as much as I can, it's becoming viscous with foaming and
bumping. Neither the chromate or dichromate will crystallize out. Something else is needed.
How do you plan on converting the calcium salt to sodium? A calcium precipitate seems like the simple fix, I have obviously collected too much soluble
matter for a direct crystallization.
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JJay
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I also ran into severe bumping when trying to boil down sodium chromate solution. I noticed much less bumping with sodium dichromate than with sodium
chromate, but I also used less heat and had less unwanted material in the sodium dichromate solution, so I'm not exactly sure what made the
difference.
I do know that you need to decompose the clorates somehow, or they will make it very hard to crystallize the material.
I think adding sodium carbonate or sodium sulfate would precipitate the calcium. The ideal procedure would add only exactly the required amount to
eliminate some purification steps later.
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Fulmen
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An acid boil should take care of the chlorate, right? But what about perchlorates? I did boil it down in alkaline conditions, so I could have some
perchlorates as well.
As for precipitating the calcium you still have to deal with the low solubility of calcium chromate. Adding solid chromate to a carbonate/sulphate
solution would coat the solids with the precipitate which could make reaction slow at best.
Edit: According to this: http://www.drugfuture.com/chemdata/calcium-dichromate-vi.htm... calcium dichromate will decompose into calcium chromate and chromium trioxide.
That would be a useful end product, right?
[Edited on 26-6-17 by Fulmen]
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JJay
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Both the calcium chromate and the carbonate or sulfate would need to be in solution before precipitating the calcium or it would likely make a big
mess.
Chromium trioxide is a useful end product, but I'm not 100% sure how you would remove it from the calcium chromate... I wonder if extracting with
acetone would work.
While calcium dichromate breaks down at only 100C, sodium chromate can withstand the temperatures needed to decompose sodium chlorate and sodium
perchlorate.
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Fulmen
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Acetone could work. Wikipedia lists it as soluble, but I haven't found any numbers. Since Cr(IV) is used to oxidize alcohols into ketones it's not
unreasonable to assume that it won't oxidize any further.
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Fulmen
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Well, at least I'm getting a precipitate now:
http://www.sciencemadness.org/talk/viewthread.php?tid=26378&... .
I merged the two extracts, acidified it and was able to reduce it down quite a bit more. Also, chlorine gas.
Hello Chlorine, my old friend
I've come to inhale you again
Nah, once bitten, twice shy. Painful as hell, so I'm not doing that again. I'm assuming it's sodium chloride but it's quite yellow so some trapped
dichromate must be recovered.
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JJay
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My crystals never seem to look as nice as everyone else's, but the sodium chloride seemed to separate cleanly when I tried that with sodium chromate.
Of course, you don't want any chloride in your dichromate product, and measures should be taken to verify that it doesn't contain any. Probably the
best way to do that is with a chromyl chloride test.
Definitely try to avoid inhaling any hexavalent chromium; it's a known human carcinogen.
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Fulmen
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I know. I actually handled it by the kilos at a lab I was interning. Scary shit, not something you want to spend a lifetime working with.
A few observations: I started working on a new batch of leached metal, this time I heated the precipitated product to a dull red before further work.
This produce a denser powder that is far easier to work with than the precipitated carbonates.
15% NaOCl does wonders for the reaction speed, but it also interferes with the settling process. Gas (presumably O2) is generated, keeping the
solution agitated. So I'm forced to filter off rather than settling/decanting.
Acid destruction of hypochlorite works fine as long as you can deal with some chlorine gas.
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JJay
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I might have to give acid destruction of hypochlorite a try.
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Fulmen
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Judging by the amount of noxious gas and salt I'm getting it's quite effective. I started acidifying it cold, once the out-gassing subsides it can be
boiled. It's also obvious that I need to let the hypochlorite react for longer as I'm clearly wasting most of it.
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Fantasma4500
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only chromium hydroxide well soluble in caustic solution? i would suppose solubility of metal hydroxides to increase with pH - if this is the case
then screw oxalate, hydrogen peroxide is not ideal for oxidation, ferrate is barely possible to use, its on par with francium in my world, ive barely
been able to produce a noticable hint of it in solution years back! (ferrate, that is.)
i would agree that using sodium chloride rather than hydrochloric acid would be the better way to go, usually no free chlorine gas or hypochlorite
even is formed as it reacts with the metals in stainless steel before anything else, i have used electrolysis to kickstart dissolution of stainless
steel with HCl however. one problem with using NaCl and electrolysis would be having a billion scrap pieces of stainless steel, you would need to tap
weld all the little things together to get a reasonable anode, could glue work? maybe using smaller stainless steel pieces as neutral electrode, in
direct contact with the anode itself, but this yet again calls for a larger piece of stainless steel, maybe smaller scrap pieces of stainless steel is
best dealt with using acid anyhow.
dealing with slimey hydroxides is easiest done by boiling the whole thing dry, actually helps if theres a lot of soluble solids in it, as this will
cause it to become a more brittle solid once somewhat dried out, once water is added again it doesnt tend to turn into a slimey hell clogging up
filters and causing a mess, and most importantly - mocking your final yields.
removing chlorates can be done quite carefully by mixing in a bit of fuel, such as sugar, and then adding a bit of acid, the HClO3 formed should at
quite low concentrations start to react with anything combustible, rule of thumb with HClO3 is that +30% and it will be capable of in solution
detonating, of my memory it reacts with sugar in very low concentrations
if i didnt mention it already, manganese in steel alloys can be a real hazard when intending to use hypochlorite as oxidizing agent, i think whats
important is to not dry out the mix, at least with precipitated chromium hydroxide from oxalate filtrate this caused a lot of chlorine to be formed -
still completely unclear on whether it could have been risidual oxalic acid that caused decomposition of the hypochlorite, as mentioned earlier in
this thread iron can apparently complex with oxalic acid.
working with nickel compounds in solution, managing this solution can give you nickel poisoning, just by carefully evaporating the water out of a
nickel containing solution
so for now.. electrolytic dissolution of stainless steel using NaCl, followed by ppt. of iron and nickel hydroxides while keeping chromium hydroxide
in solution using high caustic pH, followed by hypochlorite oxidation (after neutralization of NaOH with HCl into NaCl) (MMO electrolysis of NaOH,
Cr(OH)x, NaCl?) of chromium hydroxide into sodium chromate
which can then be fractionally crystallized out or more easily ppt by turning the sodium chromate into potassium chromate, acidified with HCl into
dichromate
i would avoid dichromate as a mean for isolation as it can form chlorochromate which by my experience can decompose into chlorine, unsure if this
decomposition happens considerably faster with heat than it does at ambient temperature.
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Diachrynic
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This pdf contains some interesting charts about iron / chromium:
Attachment: Potential - pH diagrams of Cr-H2O system at elevated temperatures.pdf (367kB)
This file has been downloaded 374 times
we apologize for the inconvenience
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Fulmen
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Decided to revive this project, on a whim I figured I could try the calcium chromate route. Big mistake, the precipitate was so fine it was impossible
to filter even with vacuum. So I ended up redissolving it in HCl and precipitate the calcium with sodium sulfate. Weirdly this produced the most
wonderful, dense precipitate that settled out like a dream, from the solubility I would have predicted the opposite.
Well, that was a lot of work for nothing. But at least I got around to making an adapter for the aspirator so I can vacuum filter again. That was the
reason for the delay, my fridge compressor died after too much abuse.
Guess potassium dichromate is the best way to go. Sadly I don't have enough potassium carbonate for this, the only other source of K I have is low
sodium salt (50% KCl). But I guess I'll have to make it work, gotta finish this now.
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