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Author: Subject: Seperating chrome from iron and nickel, stainless steel
JJay
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[*] posted on 23-5-2017 at 01:50


After drying, there are 59.20 grams in the first fraction. The second fraction (from evaporating 500 mL of the volume) is more voluminous and was also much easier to filter. It has a sandy, sparkly, crystalline texture and is hopefully pretty pure.



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[*] posted on 23-5-2017 at 07:32


actually it seems calcium chromate has upside down solubility curve: https://en.wikipedia.org/wiki/Solubility_table
meaning you want to filter the solution hot as thats where solubility is minimum

the reason i dont favor dichromate is simply because im talking about the purification part, dont wanna be messing around with toxic filter papers when i can avoid doing so, turning the chromate into dichromate is dangerously simple, its a seperate part though

actually calcium chromates solubility curve is quite desirable, as a reversed solubility curve is rare this means that you can get most other stuff out by dissolving it in cold water, to then re-precipitate most of it from a hot solution

your procedure for this however.... was it that you reacted stainless steel with Ca(ClO)2? (what about electrolysis of SS in CaCl2?)
i can only imagine how much of a hassle that must be, whenever i have chance of getting straight to precipitation i will do so because decantation is doable at a very large scale, and you barely need to deal with paper or filters or whatever, i also have an idea that strapping a vibrator to a container with liquid and solids that you want to have seperating would greatly aid the solids in settling

promising looking yield, one way to tell a precipitate being quite free from soluble matter is whether it falls nicely apart when dried up, if soluble matter is still in the solid as it dried out it would suggest crystallization took place




~25 drops = 1mL @dH2O viscocity - STP
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[*] posted on 23-5-2017 at 08:34


Quote: Originally posted by Antiswat  
actually it seems calcium chromate has upside down solubility curve: https://en.wikipedia.org/wiki/Solubility_table
meaning you want to filter the solution hot as thats where solubility is minimum


That seems to contradict other information that I have on the subject, not to mention the main calcium chromate Wikipedia page. The calcium chromate solubility curve is not nearly as steep as, for example, potassium dichromate's, but I don't think it is inverted.

I haven't weighed the second fraction yet, but it is looking pretty nice. Right now, the third fraction is filtering. I could smell chlorine and chlorine oxides being given off as the last bits of its 500 mL solution evaporated, and I can see a little bit of very fine white precipitate in it the filter cake. That actually surprises me because I hadn't smelled any chlorine previously. I think the smell indicates some calcium hypochlorite was present and might be in the crystals, but I don't think they contain much of it.

Aside from calcium hydroxide, most of the likely impurities, i.e. calcium chloride, calcium chlorate, and calcium chlorochromate are deliquescent. There could very well be around 350 grams of calcium chloride left in solution. Saturated at room temperature, it could take up about 550 mL of solution. The next two 500 mL fractions might contain some product but will be very impure, especially the final one.




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[*] posted on 23-5-2017 at 11:19


The second fraction is 147.4 grams, which is quite a lot for only 500 mL of solution. It is a little bit clumpy, so I don't think it is quite 100% dry, but it's not deliquescent. It should be recrystallized for purity and dried in a dessicator I think. The third fraction was harder to filter (though not as hard as the first) and is much fluffier. It has some variation in color, with a pale substance mixed in with the much brighter yellow powder, and it looks smaller than the second fraction. It will definitely need to be recrystallized.



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[*] posted on 23-5-2017 at 17:29


After boiling off another 500 mL, I can see only a little bit of precipitate, and a syrupy yellow liquid remains. This is in fact a clear solution with a little bit of yellow precipitate and some suspended yellow particles. It is very hard to filter.

IMG_20170523_171410[1].jpg - 435kB




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[*] posted on 24-5-2017 at 03:06


Fraction 3 weighs almost exactly 106 grams.

I managed to filter a small amount of some deliquescent yellow goo out for fraction #4. It's hard to really say how much, but there might be 30 grams of extremely impure powder here after drying the yellow mustard-like filter cake (which I'm guessing might be possible in a vacuum over sodium hydroxide). The filter paper has been turned to gel and can't be easily removed. The syrupy filtrate isn't 100% clear but is actually slightly yellow and likely contains very little of the desired product.

It looks like I have 312 grams of powder here. Assuming that all of the utensils I dissolved were 18/8 steel, the theoretical yield is actually only 244 grams. I know that a couple of them were chrome plated, so it's conceivable that there could be a slightly higher yield than the theoretical, but likely the inflated yield is due to impurities. I can see that the earlier fractions are brighter yellow than the later fractions. Clearly, a recrystallization is called for.

IMG_20170524_025326[1].jpg - 441kB

[Edited on 24-5-2017 by JJay]




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[*] posted on 24-5-2017 at 15:10


I stirred all four fractions in about 2L of hot water for a couple of hours and then filtered out the insolubles, which included some pieces of filter paper and a little undissolved powder (likely other chromate salts and calcium hydroxide). The filtration was easy, and I could see the lemon yellow solution take on a visibly deeper shade of yellow after it was filtered. Thinking that long exposure to filter paper might have reduced some of the chromate, I then added a little hydrogen peroxide. A very slight darkening was observed, but it passed within less than a second. Whatever was causing the solution to take on an opaque brown appearance that lasted for several minutes with peroxide (perhaps manganese or iron salts) seems to have been largely eliminated (although if I do this again, I'll try a more scientific approach by measuring the peroxide and attempting to quantify the color changes).

[Edited on 24-5-2017 by JJay]




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[*] posted on 26-5-2017 at 06:24


Here is a different path that is easy on the reagents employed and relatively simply, but slow. It is original as it is based on an observation of what one is not suppose to do when performing an electrolysis of a concentrated MgSO4 solution.

The warning is that one should not use stainless steel electrodes when performing the electrolysis as they are claimed to be attacked by the forming H2SO4 and perhaps also some sulfate radicals (namely, .SO4- , and, to a much more limited extent possibly, .SO4- + .SO4- = S2O8(2-), putting yet more species into the mix).

Corrosion of the the steel electrodes implies with certainty the formation of sulfates, and depending on conditions, along with the expected Mg(OH)2, perhaps some basic sulfates. One may decide not to stir the solution and limit air contact to hopefully reduce basic sulfate creation, however their creation, from stirring and at the end of the electrolysis, further addition of base to a 3.8 pH, may provide a known separation process. For example, see "Preparation of basic chromium sulphate from iron-chromium alloys", US Patent 2766101, link: https://www.google.com/patents/US2766101 .To quote:

"It is an object of the present invention to employ a precipitation method to produce an effective separation of chromium from iron-chromium alloys containing usual impurities such as carbon, silicon, cobalt, nickel and copper.
Another object is to obtain a granular basic chromium sulfate precipitate, substantially free from iron and other impurities, which can be readily converted to other chromium chemicals.
Other advantages and aims of the invention will be apparent from the following description.
In accordance with the present invention an ironohromium alloy is digested with a dilute sulfuric acid, at a temperature of between 60 and 105 C. The resulting pulp is filtered to remove the residue which consists largely of silica. The filtrate is added, under substantially non-oxidizing atmospheric conditions, to a hot aqueous solution of an alkali or an alkaline earth metal base, for example sodium carbonate, to a final pH below 3.8, to yield a granular basic chromium sulfate precipitate which is easily filtered and washed free of soluble salts. The temperature of said alkali or alkaline earth metal base should be above 60 C. during this step in order to insure the obtaining of a granular precipitate, since at less than 60 C. the precipitate becomes gelati nous.
The basic chromium sulfate may be heated to approximately 1000 C. to produce a high grade chromic oxide or it may be readily dissolved with acid to yield chromium solutions for producing chromium chemicals of high purity.
As the filtrate is added to the alkali or alkaline earth metal base an initial precipitation of iron and chromium occurs at a pH of about 9.0. Then as additional filtrate is added and the pH reaches about 6.0 the initial iron precipitate goes into solution.
With further addition of filtrate the complete resolution of the initial iron precipitate is brought about leaving exclusively a precipitate of granular basic chromium sulfate."

Here are also some comments per Wikipedia (link: https://en.m.wikipedia.org/wiki/Chromium(III)_sulfate) to quote:

"Chromium(III) sulfate usually refers to the inorganic compounds with the formula Cr2(SO4)3.x(H2O), where x can range from 0 to 18. Additionally, ill-defined but commercially important "basic chromium sulfates" are known. These salts are usually either violet or green solids that are soluble in water."

The good news is that Cr2(SO4)3 is around 5 times more soluble than Iron and Nickel sulfates, which could be of value in fractional crystallization, assuming one does not select a basic sulfate processing route outlined in the cited patent.

[Edited on 26-5-2017 by AJKOER]

[Edited on 27-5-2017 by AJKOER]
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JJay
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[*] posted on 26-5-2017 at 06:36


The major problem I've been running into is heating too strongly, which causes a crust of hard-to-dissolve material to form in the bottom of the beaker. I manged to get most of it dissolved by adding a couple liters of extra water and then evaporating it off slowly. It's filtering right now.

I like the idea of using electrolytic methods to purify the metals.




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[*] posted on 1-6-2017 at 01:30


working with ppt of iron-nickel oxalate yesterday as well as managing the solution of chromium oxalate gave me slight tremors, supposedly from addition of sodium carbonate which formed lots of bubbling and thus knocked a bunch of the solution into air, i misjudged pH quite a lot it seems as stainless steel pan i poured the chromium oxalate and oxalic acid solution into had been darkened a great bit and the green solution now keeps giving off iron, i know it was going too easy. tremors were minimal and gone in the morning however, one can take it as an experience maybe.

so knowing oxalic is quite strong, how far out is it to let oxalic acid itself devour stainless steel? a concentrated hot solution of oxalic acid seems to give as nightmarishly irritating..... well i wouldnt say fumes because it doesnt seem to fume, but irritation similar to sulfur trioxide, if you can imagine inhaling nails, so you def want to avoid oxalic acid solutions going too hot without strong ventilation or likewise, adding sodium carbonate solution im now getting lots of very nice lightblue precipitate, supposedly quite pure chromium carbonate, completely ignoring iron-nickel that dissolved from the pot itself, the iron however over time comes out of solution as iron oxide which is easily skimmed off, worryingly easily.

a route to hypochlorite would be adding a strong base to TCCA, tricholoroisocyanuric acid which is sold as solid pool chlorine, the hypochlorite formed could be used effectively in situ, woelen has had quite a craze over TCCA and indeed it has a lot of use still to be found

as a bonus with oxalate process you can use iron oxalate to form iron nanopowder, which mixed up with NaNO3 could reduce the NaNO3 into NaNO2, this is almost theoretical, supposedly the nickel metal would be left unharmed by molten or at least hot NaNO3/NaNO2
the chromium sulfate process seems quite technical but in large amounts with great control very likable

on a sidenote.. about pigments; once had a small bag of lead chromate, i found it to be actually PbO*PbCrO4 which was used as a pigment back in time, a very strong orange colour, very insoluble, not sure how difficult it would be to make the PbO*PbCrO4 rather than just PbCrO4, i suggest you could maybe ppt out the remaining chromate using barium chloride (which is soluble itself) then you would be dealing with barium chlorate which is also soluble, barium chlorochromate would be up to you to find out about

chromic acid could be a thing for purification even, it seems somehow unstable but distilling it over at some 200*C

im gonna have a go at oxalic acid + stainless steel, this could skip the first acid dissolution/conversion step, i would however advice to use NaOH in first go to avoid using heaps of Na2CO3, knocking heavymetals into the air and you, a bunch easier etc, although being respectful to chromium carbonates solubility in alkaline solution




~25 drops = 1mL @dH2O viscocity - STP
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https://en.wikipedia.org/wiki/Solubility_table
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[*] posted on 1-6-2017 at 08:21


I need to transport my lab over here to finish this, but at this point I have a bunch of crystals and about 500 mL of liquid remaining.



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[*] posted on 2-6-2017 at 21:01


I got my lab transported. First thing I am going to do is to finish working up my calcium chromate....



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[*] posted on 3-6-2017 at 03:49


when drying out chemicals, forcing it through a sieve with a dowel makes the drying procedure quite a lot easier

kinda feel like i have to point out that oxalate seperation of chromium seemingly makes it so that there is also manganese in the chromium carbonate mixture, chlorine gas concentration paired with its immense density can really take you by surprise, the shock you get from chlorine is probably exponential with concentration

i cannot quite make sense out of why the hell there would be a bunch of manganese left in the chromium carbonate as manganese oxalate should have found its way out of the equation along with the iron and nickel oxalate, effectively decomposing about half a litre of bleach just doesnt smell like trace contamination




~25 drops = 1mL @dH2O viscocity - STP
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https://en.wikipedia.org/wiki/Solubility_table
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[*] posted on 21-6-2017 at 09:55


Antiswat: Can you confirm the solubility of Cr(III)Oxalate? I could only find data for Cr(II)Oxalate.



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[*] posted on 21-6-2017 at 12:32


I found in doing my workup that repeated fractional crystallization does improve the purity of calcium chromate, but I think it is probably more efficient to purify it through chemical means, such as reduction and removal of the chromium by filtration followed by re-oxidation with a clean oxidizer like hydrogen peroxide or perhaps decomposition of any residual chlorates at high temperature followed by conversion to potassium dichromate, which is much easier to recrystallize. If you want to get really fancy, you could convert chromates to chromyl chloride, distill it, reduce the chromyl chloride, purify through filtration, and re-oxidize.



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[*] posted on 22-6-2017 at 06:04


JJay: Did you get any black precipitate after treating with OCl? I'm using sodium hypochlorite on the carbonates, and the filtrate started out a clear yellow. After some boiling I got a black ppt while the solution has turned more greenish, I'm suspecting manganese oxidized to permanganate before decomposing to MnO2.



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[*] posted on 22-6-2017 at 06:24


I didn't notice any black precipitate forming after I filtered. I did notice a transient black color when adding hydrogen peroxide... I suspected manganese but wasn't 100% sure about it. Is it possible that your chromate is getting reduced somehow?



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[*] posted on 22-6-2017 at 08:05


Don't see how, as it's a clear solution in glass. Heat alone shouldn't be enough, right? I actually got some ppt before boiling after the filtered solution was left alone over night.

Edit: Manganese fits perfectly, doesn't it?
It can form manganates in the presence of hypochlorite which should decompose with heat. The precipitate (not much) was easy to filter off, although some stuck to the glassware like a coating. So it's a simple step that should remove most of the manganese.

[Edited on 22-6-17 by Fulmen]




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[*] posted on 22-6-2017 at 21:43


Today I tried putting some hydrogen peroxide into some technical sodium dichromate solution. It also produced a transient black color. I'm starting to wonder if maybe it is forming a higher oxidation state of chromium or (more likely) actually reducing it and then re-oxidizing it. I've seen the black color form with peroxide from sodium chromate, potassium dichromate, sodium dichromate, and calcium chromate now, each from different sources, so if it's caused by an impurity, it's extremely widespread. The duration of the color seems to vary considerably, though, and I'm not exactly sure what factors affect it.

Distilling chromyl chloride would eliminate manganese contamination... surely there is an easier/safer way, though....




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[*] posted on 22-6-2017 at 22:18


Looks like the transient black color I've been seeing may be this: http://www.sciencemadness.org/smwiki/index.php/Chromium(VI)_oxide_peroxide

I'm not yet sure how to ascribe the color changes I've seen it take, but I imagine that it is affected by the amount of peroxide, heat, concentration, stirring, and other components in solution.

Apparently its complexes are actually pretty useful: http://www.sciencedirect.com/science/article/pii/S0040402001...

Reading over the paper, they also look easy to prepare, although a lot of ether is required.


[Edited on 23-6-2017 by JJay]




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[*] posted on 22-6-2017 at 23:58


Sounds like a reasonable assumption. It's not the same as I've been seeing, I am getting a solid, stable precipitate.



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[*] posted on 23-6-2017 at 03:24


I don't know for sure what that is... you could test the precipitate using a procedure from this paper: http://pubs.acs.org/doi/abs/10.1021/ed060p134

Quote:

Oxidation to Permanganate in Basic Solution Test: Mn(II) test solutions as low as 0.001 M can he used in 1-2-drop portions. These test solutions can be prepared from the MnO2 produced in the qual separation by dissolving it in a minimum of H202 and HCl and boiling. 1) Prepare the oxidant: Mix 10 drops of saturated NaHCO3, 1 drop 0.05 M CuSO4, and 10 drops 6% NaOCl (bleach). 2) Add to the oxidant 1 drop of Mn(II) test solution, about 0.01 M. Heat in a boiling water bath for 2 min. along with a blank. If the test color is pale, add more of the test solution, and continue heating. Chloride, present in NaOCl, does not interfere. This method was suggested by Feigl (2); however, much precipitate (CuO, MnO2?) form when NaOH is used alone. Following a hint in Mellor (3), we found that no precipitate forms when a bicarbonate buffer and less Cu(II) are used. A possible explanation of this result follows. Cu(II) and Mn(II) are held in solution as carbonate-bicarbonate complexes. OCl- produces a Cu(III) species which also remains complexed in this mixture (1). The Cu(III) or some other transition metal catalyst is needed for some of the oxidation steps (possibly the final MnO4 2- to MnO4- reaction), since OCl- does not produce appreciable MnO4 by itself. We found that a pure green MnO4 2- solution in 0.2 M NaOH was oxidized by the Cu(III) oxidant in 2-3 min at room temperature. This test in basic media is useful because, in the presence of much chloride, the usual acidic oxidations (by persulfate, periodate, or bismuthate) require excess Ag+ or Hg+ to prevent chloride reduction of permanganate which is rapid in acid (4).






[Edited on 23-6-2017 by JJay]




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[*] posted on 23-6-2017 at 03:59


Your comment on Iron oxalate solubility is true and untrue, as it readily complexes with excess H2C2O4 to again be soluble!

In fact, commercially sold iron stain removers use Oxalic acid!
---------------------------------------------------------------

Take another look at my provided patent reference above, it may save you some headaches.
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[*] posted on 23-6-2017 at 12:34


Nah, I don't think it's that important. Whatever it is, it shouldn't be chromium so good riddance.

I am surprised how slow this reaction seems to be. I use 2.4% NaOCl, and the washings just keep on coming out yellow. The solution still smells of hypo after several hours of stirring, and yet the next wash looks almost the same. Problem is that I tried filtering the sludge rather than just decant it, and the hypochlorite did a number on the filter paper. This dissolved paper seems to act as a thickener, making further settling impossible.

Guess I'll have to reduce it to dryness, burn off the paper and repeat. Anything I can do to improve the yields?

Did you choose calcium hypochlorite for it's availability/price or the solubility? To me it was both, I find soluble compounds much easier to work with. Precipitation products often end up as fine, unworkable powders. After messing about with copper chemistry I developed a distaste for this, so I will always choose crystallization of soluble salts if I have a choice. They provide infinite do-overs.





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[*] posted on 23-6-2017 at 12:53


I chose calcium hypochlorite for the price mainly, and also, I didn't want to use several jugs of bleach. One problem is that the iron oxidizes along with the chromium, wasting oxidizer... if you have to process a lot of stainless steel, it might actually be a good idea to look at the basic chromium sulfate patent that AJKOER provided above.... I let the calcium hypochlorite react for weeks, and the hypochlorite smell never went away entirely. I think you can actually use a much shorter reaction time. I also ran into that same problem with the filter paper weakening and turning to gel.

I'm not really sure how to improve yields... People have attempted to conserve oxidizer by using air to oxidize the iron (such as in this video: https://www.youtube.com/watch?v=mjjIzie_E3s). But it's not as easy in practice as it is in theory.

I'm pretty sure the easiest way to obtain high purity is to decompose the chlorates and convert to potassium dichromate then recrystallize.




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