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Fantasma4500
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Seperating chrome from iron and nickel, stainless steel
been curious about dragging chrome out of stainless steel because its quite abundant and the element chrome for whatever reason seems interesting to
me, fancy colours or whatever
anyhow the common procedure for me is to dissolve it in HCl, thereafter NaHCO3 it into basic chromium hydroxide, filter that off (along with lots of
iron and nickel hydroxide)
react that with NaClO, then NaCrO4 is produced, it can at this point then be fractionally recrystallized, possibly crystallizing it as large crystals
hoping it wont co-crystallize with other nasties in solution
anyhow.. it seems that chromium oxalate is quite soluble, more than 100g/100mL, i dont get why it would be, it seems like it doesnt quite follow the
system metal oxalates usually do, where nickel and iron are quite insoluble
Nickel oxalate NiC2O4.2H2O 0.00118
Iron oxalate dihydrate 0.097 g/100ml (25 °C)
Chromium oxalate 126 g/100 mL (0 °C)
to eliminate the excess oxalic acid that one would intentionally dump into the dissolved stainless steel, a solution of iron chloride would be added
just until no more ppt forms, maybe a few drops excess iron chloride, the iron chloride would then with air react to ppt the iron chloride, leaving
despite trace metals in stainless steel (manganese 0-2%, molybdenum, 316 SS?) a quite pure solution of chromium oxalate
i had previously thought about decomposing the whole lot of oxalates into corresponding oxides/hydroxides
iron oxalate as many know decomposes into nano iron metal, it then quickly reacts with air and moisture, it could possibly upon formation be removed
per magnet assuming it would be fine powder, it does tend to form oxides and hydroxides upon decomposition however, not entirely pure iron powder.
nickel oxalate would decompose into quite pure nickel metal with trace amounts of nickel carbonate and hydroxide
chromium oxalate would decompose into Cr2O3 with CrO and Cr3O4 as intermediates (or the other way around) anyhow through decomposition of the metal
oxalates it could be possible to dissolve the chromium oxides in strong alkali solution
now as i see it there is a chance of even dragging nickel and chromium somewhat unharmed out of stainless steel solution simply adding zinc, the iron
would naturally through air exposure turn into hydroxides and oxide, a weak acid could be used to leech out the iron hydroxide, where Fe2O3 of what
ive tried seems quite resistant to even HCl, nickel and chrome should dissolve a lot faster
"The highest chromium leaching was achieved with the aqueous solution of oxalic acid, as chromium was converted into water-soluble chromium oxalate."
http://agris.fao.org/agris-search/search.do?recordID=US20130...
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JJay
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I'm going to work up some chromates that I produced in pretty much the same way with Ca(OCl)2 as soon as my new filter funnel arrives. One thing I
have wondered about is whether it might be possible to leach chromium out of dissolved stainless steel using sodium hydroxide to make
Na3Cr(OH)6.
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Fantasma4500
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i doubt that, 316 SS is resistant to boiling caustic soda, it tears only few millimetres per years exposure
electrolysis may be a thing, but im quite sure iron and nickel hydroxides will also be soluble
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JJay
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The procedure would involve first dissolving the stainless steel in hydrochloric acid and then neutralizing with sodium carbonate.
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Fantasma4500
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thing is you get a mix of iron nickel and chromium carbonate/hydroxide
purity is the problem at that
you may try fractionally recrystallizing the NaCrO4 however
another way to get iron out of the mixed stainless chloride solution would be letting oxygen get at it, sadly it seems hydrogen peroxide wouldnt aid
this reaction much as it would form actually more stable FeCl3, airbubbling the thing would be possible but still extremely messy and taking long time
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JJay
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Nickel hydroxide is not soluble. Iron hydroxide is, but iron oxide is not. It could be tricky to remove all traces of iron that way; it would be
easier to simply crystallize potassium dichromate, but with sodium hydroxide as the limiting reactant, I don't know if that would even be necessary.
No matter how you do it, it's going to be messy, but I think extracting with sodium hydroxide looks like a promising alternative that hasn't been
tried by amateurs, and economy of reagents looks good that way.
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Melgar
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Quote: Originally posted by JJay  | | Nickel hydroxide is not soluble. Iron hydroxide is, but iron oxide is not. It could be tricky to remove all traces of iron that way; it would be
easier to simply crystallize potassium dichromate, but with sodium hydroxide as the limiting reactant, I don't know if that would even be necessary.
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I'm pretty sure that in the presence of oxygen, iron hydroxide forms iron oxide-hydroxide, which is that reddish-brown rust we all know and love. And
that's quite insoluble in water. It might actually be easier than you think.
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JJay
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I think you're right, and that process can be accelerated considerably by using hydrogen peroxide, which would also probably form insoluble iron (iii)
oxide. I'm pretty sure there's also a risk of forming sodium ferrate this way, though.... It decomposes pretty easily, but I'm not 100% certain what
the decomposition products are... iron (ii) hydroxide is the one to avoid... pretty sure iron (iii) hydroxide is insoluble.
[Edited on 15-5-2017 by JJay]
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Melgar
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| Quote: | | I think you're right, and that process can be accelerated considerably by using hydrogen peroxide, which would also probably form insoluble iron (iii)
oxide. I'm pretty sure there's also a risk of forming sodium ferrate this way, though.... It decomposes pretty easily, but I'm not 100% certain what
the decomposition products are... iron (ii) hydroxide is the one to avoid... pretty sure iron (iii) hydroxide is insoluble. |
Let's put it this way: it's much more likely that you form an iron iii compound when trying to get an iron ii compound or especially a ferrate than
vice versa. If you did somehow get some in solution, it'd probably oxidize your chromium ions to hexavalent chromium, which would at least be very
easy to notice.
I actually had a strip of spring steel that I'd dissolved in anoxic HCl some time ago, and I finally poured it out in a dish. The solution was a
dull, dark blue-green color. I then added a solution of saturated sodium hydroxide, and if formed more of a sludge than a precipitate. I stirred it
up, but nothing really settled. The color was more or less the same though. I noticed brown forming around the edges of the liquid, which was
clearly iron oxide-hydroxide, so I dripped some 30% H2O2 in, and where each drop landed, a large brown spot formed. I then added quite a bit more,
and stirred it up. A lot more of it turned brown, and that started to settle. The liquid above the settling brown rust precipitate was clear though,
I'm not sure why that would be the case, and I can't test it until I get home from work.
edit: But I think it may be one of those dreaded double salts forming. Like chromite:
https://en.wikipedia.org/wiki/Chromite
It should at least be possible to concentrate in that state though.
[Edited on 5/15/17 by Melgar]
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JJay
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The mixture containing the dissolved silverware that I have been playing with for a few weeks (so far it's been dissolved in HCl, precipitated out
with sodium bicarbonate, washed, and reacted with bleaching powder) started to settle; it is still bubbling and smells like chlorine oxides, but the
gas being emitted is definitely less noxious than before. The mixture will still need to be filtered, but I can see a dark orange solution above a
dark brown/orange precipitate.
I'm pretty sure I'll get to it before this weekend.
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Melgar
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Those orange-colored ones are usually nothing good: a combination of yellow iron chloride and reddish-brown oxide-hydroxide in various proportions, I
bet.
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JJay
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I'm sure the precipitate has nothing in it that I really want except for perhaps some nickel oxide. I think it's probably mostly iron oxide. Hopefully
the vast majority of the chromium has leached out; my plan is to wash the precipitate until the washings turn clear and then put it in a trash bag,
but if necessary I can add some hydrochloric acid and isopropyl alcohol and let it stand overnight (preferably outside). Then I'll just boil the
washings / filtrate down in my 5L and collect crystal fractions before I decide what to do next. I'm tentatively leaning towards decomposing the
calcium chlorates in a crucible and then making potassium dichromate, but I'd definitely prefer to have sodium dichromate....
[Edited on 16-5-2017 by JJay]
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battoussai114
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I think it was the materialsproject.org database that had pourbaix diagrams for multiple metal ions. You could use it to try and find conditions for
precipitating one while leaving the other in solution.
[Edited on 15-5-2017 by battoussai114]
Batoussai.
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Fantasma4500
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hm hm.. nickel hydroxide should be possible to extract using ammonia, turning it into a complex, where iron shouldnt be very soluble with ammonia,
essentially the solution could be evaporated off to take out the ammonia, until solution is almost entirely neutral which would make it unlikely for
iron or chromium hydroxide to dissolve as they dont form water soluble compounds with ammonia, or are in general water soluble, only worry would be
hydroxides gelling up
iron hydroxide regarded as soluble, isnt that only in strongly alkaline solutions? im seeing people suggesting its only "soluble" in acidic solutions,
or in other words it isnt but its a base that can react with acids to form soluble compounds
i think its best avoided to deal with dichromate until you really want it, excess HCl can also form chlorochromate which over time and probably upon
heating decomposes the HCl and chlorochromate into chlorine gas, very neat to have a bottle of, chlorine gas can be fun to play around with
on a sidenote the impure chromium hydroxide sludge can be added to electrolysis cell, a bit of iron will come out but for whatever uses chlorate has
its merely benifitial to be contaminated with iron oxide
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JJay
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My plan was to dissolve it in hydrochloric acid and crystallize it then heat it in a crucible to destroy any organics.
How could you not want dichromate?? It's extremely useful and interesting, not to mention recyclable. The downsides are a) it's toxic and b) it causes
cancer, but I could say the same thing about 70% of my other chemicals.
On the other hand, I'm not really looking forward to filtering and boiling down 12L of calcium chromate solution to get 200 grams of crystals... but I
am going to get started with that now....
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JJay
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![IMG_20170520_013448[1].jpg - 447kB](https://www.sciencemadness.org/whisper/files.php?pid=483624&aid=59262)
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JJay
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So anyway, after many hours of boiling, washing, and filtration, I have about 4L of orange solution with some light precipitate starting to form in
it. I'm not too sure about what that is, but I suspect it is calcium hydroxide or perhaps a basic calcium salt. There are also pine needles, metal
oxides, and other undesirable solids floating around since a lot of the mixture was just decanted and not filtered. The washings from the metal sludge
are now yellow. The theoretical yield is 244 grams, which would require about 1.5 liters of water to dissolve fully, but there is likely a lot of
dissolved calcium chloride and some calcium chlorate, which decrease the solubility of calcium chromate, although to what degree I do not know. My
plan now is to continue washing chromates out of the sludge with water until I get bored doing that and then try to filter the sludge and neutralize
any remaining oxidizers with sodium bisulfite. Then I'll evaporate the filtrate down to oh I dunno... maybe 2.5L, filter it, and then start collecting
crystal fractions maybe every 500 mL or so, cooling the beaker to room temperature in an ice bucket before filtering fractions. Either that or I'll
just evaporate off all the water and put the experiment on the shelf since this takes a lot of time and produces a common reagent, although it is fun
IMHO.
[Edited on 21-5-2017 by JJay]
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JJay
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After doing some filtering and whatnot, I now have about 2.5L of orange solution. I thought it looked a little bit dingy, so I added a little hydrogen
peroxide to see what would happen. The solution immediately turned completely brown/black and started fizzing. The color was reminiscent of a
manganese compound. Over time, it gradually returned to a brighter orange.
I'm a little unsure about what the black color could have been, but I have seen it before with hydrogen peroxide when trying to recycle chromium waste
(which should have contained few mineral impurities, but it's hard to say 100% for sure). At this point, I'm not certain that this isn't some
impurity. I can't 100% for sure rule out the presence of manganese chloride, but I think it is unlikely that any is present.
I'm picking up some lab grade sodium chromate later and will see if it has a similar reaction with hydrogen peroxide.
![IMG_20170521_225357[1].jpg - 391kB](https://www.sciencemadness.org/whisper/files.php?pid=483793&aid=59280)
[Edited on 22-5-2017 by JJay]
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JJay
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The first crystal fraction (everthing above 2000 mL) contains a significant amount of yellow powder and some flat square-shaped crystals. The yellow
powder I think formed as the water boiled off and the crystals I think formed as the solution cooled. As is typical of mixtures containing calcium
salts, the mixture is hard to filter. I don't know for sure yet how pure it is, but indications are that the filter cake consists largely of calcium
chromate.
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Melgar
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Lots of transition metal ions will catalyze the decomposition of H2O2, iron being one of them. Are you familiar with Fenton's reagent and Fenton
chemistry?
Anyway, yellow is far more likely to be an iron iii salt I think. They're very persistent. And very yellow.
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JJay
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I don't see why iron iii would be crystallizing out at this point unless there happened to be an extremely large amount of it in solution, which is
extremely unlikely, and it makes sense for calcium chromate to be crystallizing about now. Actually, I think the suggestion that the precipitate is
iron(iii) is preposterous, and if you're making outlandish suggestions for pedagogical reasons, please desist - that's not actually good pedagogy. But
traces of iron compounds (likely present) might turn black with peroxide....
Obvious checks include testing its solubility both before and after heating at high temperatures, seeing if a solution produces a precipitate and a
red color with sulfuric acid, etc. But I'm going to collect all of the fractions before I do that.
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JJay
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I have very little information on manganous chromate, but this dialog suggests that it is sparingly soluble in water. Its presence in chromate salts
would almost certainly be undesireable: https://books.google.com/books?id=Nxw_AQAAMAAJ&pg=PA94&a...
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Melgar
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Just saying that from my own experience trying to separate nickel from stainless steel, getting rid of the iron salts was very difficult. It didn't
help that chromium can also be green, and mixed iron salts can be sort of green especially with ammonium present.
Also, white crystals forming in a strongly colored solution will often take on some color from the solution and need additional recrystallization to
purify completely.
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JJay
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There could be and likely is some calcium hydroxide precipitating, but keep in mind that calcium chromate is a very yellow pigment sometimes referred
to as Yellow 33; while it's been largely phased out due to both toxicity and in favor of more chemically resistant pigments like lead chromate and
cadmium chromate, it is still used in oil paints and wood stains. I'm not personally planning on using it as a pigment, but I do know an artist who
wants some and thinks it is super cool that it is made out of silverware...
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JJay
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Fraction 1. It is extremely bright yellow and sparkly with the consistency of fine sand, and given the conditions under which it was produced, I am
sure that it is substantially calcium chromate. I can tell it is not quite pure... it contains some pale amorphous material and square crystals, and
the square crystals are actually white. I'm not 100% sure what the impurities are, but my guess would be calcium hydroxide, which might explain why
this was so hard to filter. I'll have to dry everything out before I weigh it, and a recrystallization is certainly called for, but it looks like
there are 50-100 grams here.
![IMG_20170522_144237[1].jpg - 500kB](https://www.sciencemadness.org/whisper/files.php?pid=483830&aid=59286)
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