Sciencemadness Discussion Board - Seperating chrome from iron and nickel, stainless steel

Seperating chrome from iron and nickel, stainless steel

Pages: 1 2

Fantasma4500 - 12-5-2017 at 23:29

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...

JJay - 13-5-2017 at 00:01

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.

Fantasma4500 - 14-5-2017 at 01:12

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

JJay - 14-5-2017 at 01:25

The procedure would involve first dissolving the stainless steel in hydrochloric acid and then neutralizing with sodium carbonate.

Fantasma4500 - 14-5-2017 at 02:14

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

JJay - 14-5-2017 at 03:41

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.

Melgar - 14-5-2017 at 17:33

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.

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.

JJay - 14-5-2017 at 19:01

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]

Melgar - 15-5-2017 at 07:55

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]

JJay - 15-5-2017 at 13:19

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.

Melgar - 15-5-2017 at 13:43

Those orange-colored ones are usually nothing good: a combination of yellow iron chloride and reddish-brown oxide-hydroxide in various proportions, I bet.

JJay - 15-5-2017 at 14:36

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]

battoussai114 - 15-5-2017 at 15:05

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]

Fantasma4500 - 20-5-2017 at 00:36

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

JJay - 20-5-2017 at 01:02

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....

JJay - 20-5-2017 at 01:37

JJay - 20-5-2017 at 22:29

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]

JJay - 22-5-2017 at 05:05

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.

[Edited on 22-5-2017 by JJay]

JJay - 22-5-2017 at 10:10

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.

Melgar - 22-5-2017 at 10:22

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.

JJay - 22-5-2017 at 10:37

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.

JJay - 22-5-2017 at 11:24

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...

Melgar - 22-5-2017 at 11:30

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.

JJay - 22-5-2017 at 12:02

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...

JJay - 22-5-2017 at 14:53

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.

JJay - 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.

Fantasma4500 - 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

JJay - 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.

JJay - 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.

JJay - 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.

JJay - 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.

[Edited on 24-5-2017 by JJay]

JJay - 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]

AJKOER - 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]

JJay - 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.

Fantasma4500 - 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

JJay - 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.

JJay - 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....

Fantasma4500 - 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

Fulmen - 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.

JJay - 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.

Fulmen - 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.

JJay - 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?

Fulmen - 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]

JJay - 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....

JJay - 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]

Fulmen - 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.

JJay - 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]

AJKOER - 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.

Fulmen - 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.

JJay - 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.

DraconicAcid - 23-6-2017 at 13:02

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.

JJay - 23-6-2017 at 13:55

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]

Fulmen - 23-6-2017 at 14:08

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?

DraconicAcid - 23-6-2017 at 14:14

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.

JJay - 23-6-2017 at 14:35

Seems reasonable.

JJay - 23-6-2017 at 15:24

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.

subskune - 24-6-2017 at 08:05

Quote:

Surely the easiest way to separate chromium from iron, nickel, and manganese is by precipitating the other metals with conc. sodium hydroxide

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.

S.C. Wack - 24-6-2017 at 09:33

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.

Fulmen - 24-6-2017 at 13:09

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.

JJay - 24-6-2017 at 22:04

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.

Fulmen - 25-6-2017 at 00:38

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.

JJay - 25-6-2017 at 01:38

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....

Fulmen - 25-6-2017 at 13:00

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.

JJay - 25-6-2017 at 16:07

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.

Fulmen - 25-6-2017 at 21:44

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]

JJay - 26-6-2017 at 00:36

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.

Fulmen - 26-6-2017 at 04:43

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.

Fulmen - 26-6-2017 at 09:22

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.

JJay - 26-6-2017 at 10:37

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.

Fulmen - 28-6-2017 at 07:12

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.

JJay - 28-6-2017 at 08:19

I might have to give acid destruction of hypochlorite a try.

Fulmen - 28-6-2017 at 09:20

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.

Fantasma4500 - 13-3-2018 at 08:53

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.

Diachrynic - 13-3-2018 at 09:31

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 432 times

Fulmen - 25-3-2018 at 15:39

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.

j_sum1 - 25-3-2018 at 20:51

You could go for ammonium dichromate.
I have just re-read this whole thread and I didn't notice anyoine mentioning it.
Acidify the calcium chromate solution to form dichromate. Precipitate by adding a solution of ammonium sulfate. Let it settle and decant the ammonium dichromate solution. Recrystallise. This is what Tdep did and it seems to me to be a good route.

I have an assortment of Cr waste to process some time soon and I am thinking of this method.

Fulmen - 25-3-2018 at 23:14

Not a bad idea. The problem is purity, without a way to filter&wash the CaCrO4 I will have a ton of sodium chloride in it. Can it be separated properly?

Fulmen - 29-3-2018 at 19:42

Before I discard the calcium route completely, any suggestions on how to produce a better precipitate?

JJay - 29-3-2018 at 20:18

I didn't have extremely serious trouble filtering after simply boiling off water to produce crystals. The slower the water evaporates, the larger the crystals will be. I do remember having serious problems with bumping towards the end, but I think they were caused mainly by other calcium salts after almost all of the calcium chromate precipitated. The filtrations with high concentrations of calcium chlorate took a while to filter and required switching filter paper regularly because the filtrate was syrupy and degraded the filter paper. I used a 1L borosilicate Buchner with a vacuum pump for most of the filtrations.

[Edited on 30-3-2018 by JJay]

Fulmen - 30-3-2018 at 06:29

That's pretty much the brute force approach. Sadly I haven't finished the buchner funnel yet, so I'm stuck with gravity filtering. But a thorough boil did improve things a bit, and if I prioritize purity over yield I should be able to make this work. It would be fun to try reacting the solid calcium chromate with sulfuric acid to produce pure chromic acid, but I fear the calcium sulfate would mess up everything.

JJay - 30-3-2018 at 10:47

I don't think mine is extremely high purity just yet... I recrystallized it 3x, but I still don't think it's quite there.

Fulmen - 30-3-2018 at 11:29

Without an actual analysis it's just guesswork, might as well guess that it's fairly pure ;-)

I did have to add a shitload of NaOH to the solution to get things to precipitate. I started by adding NaOH until the solution turned yellow, that should indicate CrO4^-2 right? Still, without a large excess I get no precipitate when adding Ca(NO3)2. And the precipitate is whiter than I expected, shouldn't it be yellow? I wonder if I'm getting a co-precipitate of some sort...

JJay - 30-3-2018 at 11:48

Calcium chromate is pretty soluble... calcium hydroxide has low but significant solubility... barium chromate is insoluble IIRC. Iron chromate and iron hydroxide are insoluble.

[Edited on 30-3-2018 by JJay]

Fulmen - 30-3-2018 at 11:56

F*ck, I forgot about Ca(OH)2. I need to rethink this a bit, I assumed calcium chromate would be the least soluble salt. Damn.
---
I feel like such an idiot. My data indicated a low solubility for calcium chromate in hot solutions, so I figured I could precipitate it out. Instead I've spent a lot of time making and filtering off calcium hydroxide. I did wonder why the filtrate kept coming out yellow though. Jeeeez...

Anyways, it's all good. I hadn't tossed out anything so I just need to acidify and boil down a lot of liquid.

[Edited on 30-3-18 by Fulmen]

S.C. Wack - 30-3-2018 at 13:36

Potassium sulfate would not be unwise then, to get something nicely crystalline. When K, Na, Ca chromates and dichromates are made from roasting chromite ore red hot in air, potassium or sodium sulfate and CaO are useful things.

The best although high energy route I suspect is dissolving SS in acid, precipitating the hydroxide sludge, and roasting it as for chromite ore. The sludge retains some of the alkali; if someone wants K dichromate it's best not to use NaOH to precipitate.

I might have tried that and failed, but I would have done it like the KOH fusion with Cr2O3 in air, which works but sucks, at least when using pottery grade Cr2O3.

Earlier I mentioned how chromite doesn't form with excess base in the presence of the Fe, but maybe it will form a soluble and H2O2-friendly chromite when the oven-dried hydroxide sludge is heated with KOH?

Fulmen - 31-3-2018 at 13:26

I think I know where it all started going wrong. I used the data from the Wikipedia Solubility Table which lists data for the anhydrous salt not the dihydrate:

anhydrous
4.5 g/100 mL (0 °C)
2.25 g/100 mL (20 °C)
1.83 g/100 mL (30 °C)
1.49 g/100 mL (40 °C)
0.83 g/100 mL (60 °C)

dihydrate
16.3 g/100mL (20 °C)
18.2 g/100mL (40 °C)

JJay - 31-3-2018 at 14:53

I have observed that calcium chromate that is heated to dryness on a hotplate is hard to dissolve without a large quantity of water. Come to think of it, I wonder if that is a good way to purify it... by first heating to dryness and dehydrating, grinding, washing with water, dissolving with a large quantity of water, concentrating, and crystallizing....

Fulmen - 31-3-2018 at 15:20

It's worth a try I guess.
I'm about done cleaning up this giant cluster f*ck. I ran out of HCl, so a large portion of the Ca(OH)2 had to be filtered. Oh well, at least I have another project cleaning that up as a bonus. Might come in handy. Even though I used recrystallized calcium nitrate I got a lot of dark brown turbid matter in the solution that really clogged up the filters, but now that's gone and I'm left with about a liter of chromate solution that I can work with again.

JJay - 31-3-2018 at 15:46

I just tested some of my calcium chromate with peroxide... I currently have four fractions, but they are not the same fractions as before. Fraction 1 turns brown, almost opaque with .8 grams of yellow, crunchy crystals in 10 mL of water with 10 mL of 29% H2O2. It's not extremely dark brown and is certainly less dark than Jones reagent with hydrogen peroxide, but it's nearly opaque in a 20 mL vial. It has a greenish cast but when held up to the light looks more reddish.

Fraction 2 is a bit damp and does not produce a solution that is as strongly yellow. It turns brown with peroxide but not as brown. Fraction 3 is also a bit damp and is even less brightly colored and produces less brown.

Fraction 4, which is dry and was actually the first fraction taken, produces an even brighter yellow and darker brown.

JJay - 31-3-2018 at 16:06

I tried the same treatment with a mL of Jones reagent, and the addition of peroxide produced considerable fizzing and a persistent dark color that appeared blue-green when diluted.

I put .8g of fraction 1 into a vial and added 10 mL of water, followed by a few drops of sulfuric acid. Addition of peroxide had the same effect as with Jones reagent except the solution was not quite as dark and did not fizz as long.

Fulmen - 1-4-2018 at 13:36

What should this test indicate?

Personally I'm loosing my frikkin' mind here. I keep getting this dark brown, muddy precipitate that instantly clogs the filters. At first I thought it was some sort of clay/powder coating residue from the fertilizer grade calcium nitrate, but it keeps appearing every time I boil down the solution further. Any idea what it can be?

JJay - 1-4-2018 at 13:49

It indicates chromate. A while back I had noticed that the yellow substance didn't react much with peroxide; it appears that acid is required for a strong reaction. I could try extracting the chromium (VI) oxide peroxide with ether or ethyl acetate; it should be blue in solution, and it forms stable complexes with some organic molecules, but it's said to be very dangerous to isolate it in pure form.

I do remember seeing something like the mud that you describe that but only when I boiled the solution down the first time and before seeing any yellow or white crystals... I did not try filtering it but decanted from it instead. Eventually, it stopped forming. I don't know what it was or if it was the same as what you are seeing... I assumed it was nickel or iron oxides and hydroxides that made it through the early decantations and filtrations somehow.

Filtering through a bed of cotton might work for removing it, but it's awfully hard to beat vacuum filtration for that sort of thing.

Fulmen - 1-4-2018 at 14:00

It doesn't look like a simple suspended solid, settling even for a day doesn't do anything to it. Besides, a solid would most likely act like a seed and get trapped in the salt as it precipitates, right? It's not something that has passed through previous filterings either, the solution has always been crystal clear before boiling. After boiling it down it has a brown turbid look.

Fulmen - 1-4-2018 at 23:44

Re: your rest, I don't really see the point. The color and pH-response should be enough to confirm the presence of chromates, considering the reactants there's really nothing else it could be.

JJay - 2-4-2018 at 03:45

Quote: Originally posted by Fulmen

Re: your rest, I don't really see the point. The color and pH-response should be enough to confirm the presence of chromates, considering the reactants there's really nothing else it could be.

That is kind of what I thought, but it was suggested that it might be an iron compound.

One of these will make your life a lot easier: https://www.ebay.com/itm/ISO-Standard-Borosilicate-1000ml-24...

It's a Yee Chen special, but it's a useful piece of equipment.

Fulmen - 2-4-2018 at 09:30

I did use vacuum, I managed to McGyver the broken funnel good enough for a couple of runs. Besides, I now have the acrylic funnel up and running: http://www.sciencemadness.org/talk/viewthread.php?tid=81589&...

Fulmen - 2-4-2018 at 13:35

The new funnel and a very coarse filter has helped a lot. Or at least postponed the problem with the unfiltrable precipitate a bit. I'm down to a real concentrate now, but that's also where things get hard because I'm a greedy motherf. by nature. I want good yields, and that means recrystallizing the unwanted salts several times. There is a lot of it and it keeps coming out yellow.

Fantasma4500 - 2-4-2018 at 23:46

for fine particulate you may attempt decantation, im however ignorant of the relative density of calcium chromate, particle size also means something in terms of how long it takes to settle, for instance silver is quite heavy, but nanoparticulate silver wont settle in a whole year.

how about lead chromate? PbO*PbCrO4 regardless of how fine the ppt would be should settle quite fast.

another thing.. for calcium chromate you could try to ppt most of the remaining in solution adding some ethanol shifting polarity

a little trick for calcium chemistry is to turn the calcium hydroxide into carbonate using CO2, from here you can then apply more concentrated CO2, dry ice. this forms CaHCO3 which is actually soluble, i managed to produce this using 5% H2CO3 a year ago or so, think i did it at room temperature even. CaHCO3 appeared somewhat stable to temperature, but once it dries out and gets some heat it should all convert into CaCO3 which is then perfectly insoluble

now as the calcium route uses pool chemicals... what about trichloroisocyanuric acid? with NaOH this forms NaClO, this would mean NaCrO4 and NaCl which would be easy to seperate, filter hot and let NaCrO4 crystallize out, problem with NaClO usually is its poor concentration and cost, you would also have some NaClO3 in solution which could be a problem when acidifying to get dichromate.
NaClO3 and NaCrO4 is not very soluble in ethanol, NaClO3 seems soluble in methanol however, where NaCrO4 is not.

maybe ppt with lead wouldnt be a really great idea supposing you got a lot of chloride in solution, many years ago i attempted to ppt copper chromate from stainless steel chlorides, the ppt was sorta gellish, and copper chromate is supposedly somewhat unstable, but it could work to avoid ppting chlorides at least, if its not unstable after all then it would be doable to boil the whole thing dry where the gelly ppt of copper chromate would lose its gel like state of precipitate and you could easily leech out the watersoluble chemicals with decantation

Fulmen - 3-4-2018 at 01:05

Decanting sort of works, but it's awfully slow and inefficient. But it's not the calcium chromate causing problems, it's something else that keeps precipitating out as I reduce the volume. Using a coarser filter does help, but it still clogs up after a while. And of course a lot of this ppt passes through, so I will have to deal with it eventually.

As for lead chromate, how would you turn that into a more soluble salt afterwards?

Fulmen - 4-4-2018 at 12:06

I feel I've come as far as I can with the calcium route. Either I'm getting chromate precipitation or the liquid is simply so concentrated the precipitated white salts look yellow. Either way it's getting hard to work with the solution now. I also realize the high pH was a mistake, it caused carbonates to precipitate onto the bottom of the beaker.

So I decided to precipitate out all the Ca with an excess H2SO4. But now I'm faced with another dilemma. My initial plan was to precipitate potassium dichromate as it is less soluble than the various sodium salts that would be remaining. By precipitation out the colored dichromates it should be easy to see when it's all out, right? But this probably won't work with sulfates as both the sodium and potassium salt has comparable solubility when cold.

Pages: 1 2