Sciencemadness Discussion Board - Sodium Chromate from stainless stell 18/10

i want to use only otc products so i koh is off limits

KOH is OTC. it is used in soap making, among other things.

weiming1998 - 27-9-2012 at 03:23

Combining the mixed chromium and iron hydroxides with calcium hypochlorite (pool chlorine) and water, heating slightly, then filtering gives a solution of very soluble calcium chloride and the much less soluble, bright yellow, solution of calcium chromate, which can be processed into potassium dichromate (or any alkali metal/ammonium chromates/dichromates) through double displacement reaction with an alkali metal/ammonium sulfate salt.

The solution can come out purple. This is due to the formation of ferrates. Just add some household ammonia until the purple disappears and filter solution. Chromates are not strong enough oxidisers to oxidise ammonia, but ferrates will oxidize it to N2. This also destroys excess hypochlorite, which will be eliminated as chloramines and a precipitate of Ca(OH)2.

Another way that is similar to this is to use liquid chlorine (concentrated NaClO(aq), not Cl2(l)!) or ordinary bleach to oxidize the chromium hydroxides directly to sodium chromate. Most impurities can be eliminated by crystallizing them out since they aren't very soluble compared to sodium chromate.

[Edited on 27-9-2012 by weiming1998]

bbartlog - 27-9-2012 at 03:39

Quote: Originally posted by triplepoint

KOH is OTC. it is used in soap making, among other things.

I don't think so. What exactly is OTC varies from place to place - but if you have to order it over the internet, it doesn't qualify. Dudadiesel or soapgoods.com no doubt have it, but if you can't get it from a local store... that's not OTC.

plante1999 - 27-9-2012 at 04:06

Don't forget that if you want dichromates salts you need to pass CO2 gas in the chromate solution and then add a satured solution of potassium chloride to your sodium chromate solution. Cool it down and filter precipitated potassium dichromate.

2Na2CrO4 + 2CO2 + H2O -) Na2Cr2O7 + 2NaHCO3

Sulphuric acid do the job too, but CO2 is cheaper and safer.

[Edited on 27-9-2012 by plante1999]

blogfast25 - 27-9-2012 at 05:12

Salmo:

What I do is really slightly different because my purpose is to separate the Cr from the Fe, not necessarily to actually prepare chromate or dichromate salts.

I oxidise the ferrous/chromic chloride soup first with strong H2O2 (35 %), then add strong alkali (NaOH or KOH) which dissolves the Cr as chromate (VI) and precipitates hydrated ferric oxide. Filtering leaves behind the Fe (and others). The yellow solution of chromate is then the starting point for whatever you want to do.

For instance, acidify the solution and add more H2O2 (or alcohol) and the chromate is reduced back to Cr(III). This can be precipitated again with Na2CO3 as Cr(OH)3.nH2O and isolated by filtering/washing. And from that chromites/chromates/dichromates can be easily prepared without any junk accompanying them.

But plate’s method may be quicker if you’re really looking to make the chromate or dichromate of potassium or ammonium.

Re. KOH: in the age of ‘Biodiesel’, unless you’re living in Timbuktu, KOH is easy to get. The streets of eBay are paved with KOH sellers, all pack sizes imaginable. So if bbart wants to use personalised definitions of OTC that’s his problem. AFAIC.

Salmo - 27-9-2012 at 10:56

weiming your method is interesting but i haven't calcium hypochlorite,i will search it in some hardwarestores maybe i could try with simple bleach i will think about it.

plante thank you for the info

you are right CO2 is much better, and cheaper too.

blogfast you are helping me a lot but please tell me if i'm right..
H2O2 oxides FeCl2 and CrCl3 to ferric oxide and chromic oxide? is it chromium VI?
NaOH reacts with ferric oxide and chromic oxide to form ferric hydroxide (precipitate) and sodium chromate?(what's the reaction?)

Maybe I understood what you mean, you mean that first H2O2 reacts with iron (II) chloride forming iron (III) chloride then you add NaOH, you do that to avoid the coprecipitation of iron and chromium a basic complex!

Now i have NaCl NaOH and sodium chromate in solution at the same time?
I like really much your way of purification, but when i will have only my almost pure Cr(OH)3 as a precipitate how could i obtain my sodium chromate back ?

I'm stupid, I answer to myself: with stechiometric NaOH and H2O2 again

[Edited on 27-9-2012 by Salmo]

[Edited on 27-9-2012 by Salmo]

[Edited on 27-9-2012 by Salmo]

bbartlog - 27-9-2012 at 14:22

...The streets of eBay are paved with KOH sellers, all pack sizes imaginable. So if bbart wants to use personalised definitions of OTC that’s his problem. AFAIC.

I am being literal. OTC = Over The Counter; a term explicitly used to distinguish things that are available from some local brick & mortar shop from those that can be had via mail order. Not only that, but in the reagents & apparatus acquisition thread there are plenty of examples of other people using the term in the same way. For example, when a thread is titled 'OTC tert-Butanol', only two of the posters apparently use the term in the way you suggest (as a synonym for 'easily ordered' I guess); the majority cleave to the older definition and discuss consumer goods that might or might not be useful. While I wouldn't go so far as to say that *you* are the one with the personalized definition, I do think you are in the minority, and IMO your usage obliterates a useful distinction.

blogfast25 - 28-9-2012 at 05:14

Salmo:

Ooops. Going back over my lab notes I found one glaring mistake in my description above: the oxidation of Cr(III) to Cr(IV) (chromate/dichromate) with H2O2 only takes place in alkaline conditions. So the last time I separated Cr from Fe (from SS) I dissolved the SS in HCl first and filtered off any acid insoluble residue. Then the ferrous/chromic chloride solution was neutralised to fairly strong alkaline conditions. Fe drops out as Fe(OH)2 and Cr(OH)3 co-precipitates with it. Cool the slurry, then add peroxide slowly (35 or 9 %, both work – bleach should also work), this oxidises the Fe to insoluble Fe(OH)3 and Cr to soluble chromate (CrO4(2-)). Simmer the slurry for a few minutes until effervescence stop, this kills any remaining H2O2 (this is important, as we will see)

Filtering yields a yellow solution containing the chromate. Acidify and you get the dichromate.

Important: in acid conditions dichromate oxidises H2O2 to O2, with the Cr being reduced back from Cr(VI) to green/blue Cr(III). So unless you’re planning to reduce the Cr, you need to eliminate the excess hydrogen peroxide from the alkaline step. But if you’re planning to reduce the Cr(VI) to (III) anyway there’s no need to eliminate the excess H2O2 from the first step.

Obtaining the Cr as acid soluble Cr(OH)3 has the advantage of being able to synth almost any simple Cr salt from it, including of course chromates/dichromates.

Bbart: literalism is boring and context is everything. AFAIC, an eBay seller has a counter too: you order, pay and get your stuff, simples. And what you can’t buy in a brick and mortar shop you’re unlikely to be able to buy on the Net.

Salmo - 28-9-2012 at 07:35

oh my god

so today i made a big mistake.. i almost completely oxidized the solution with H2O2 only,I added something like 200 ml and it goes on fizzing, have i to throw anything away ? Is possible that with hydrogen peroxide i oxidized only the iron to iron oxide (maybe nickel too) and that the chromium is still chromium (III) chloride? What can I do now? Could I filter off the iron oxide? and than add NaOH and repeat the oxidation with the new almost iron free solution??

blogfast25 - 28-9-2012 at 08:49

Salmo:

No, don’t throw anything away. You’ve oxidised the Fe(II) to Fe(OH)3 which has precipitated (your solution was quite concentrated, not acidic enough and that caused the Fe to drop out as hydroxide). Your Cr has in all likelihood co-precipitated as Cr(III) (it tends to do that).

Now neutralise the slurry to about pH 10 – 11 and the chill it to remove neutralisation heat. If there is excess H2O2 present that might be enough to oxidise the Cr to chromate but to be on the safe side add some more H2O2, SLOWLY. The fizzing you observed before is oxygen and it means you’re wasting H2O2. In alkaline conditions the half reaction (reduction) of H2O2 is:

H2O2(aq) + 2 e- == > 2 OH- (aq), so no O2 in sight.

Then try and filter the slurry (depending on thickness, you might want to add some water first). The filtrate should contain the Cr as yellow chromate. You can test for it with a lead (II) salt: yellow PbCrO4 should precipitate.

Salmo - 28-9-2012 at 10:21

Thank you blogfast, maybe you're right but it's strange i mean: If my Fe(II) has been oxidized to Fe(OH)3 where are my chloride ions Cl-? anyway maybe It's just Fe(III) chloride but if so why is my solution so thick? Is some kind of peptization possible?
Ok I will follow you and I wont throw away anything

I will alkalize first than add h2o2, then filtrate to get the yellow solution, than (as you told me before) i will add acid (h2so4 is ok?) more h2o2 to get Cr(III) back, sodium carbonate to precipitate it and than i will wash it really well!
Then, i will have almost pure Cr(III) hydroxide as [Cr(H20)3(OH)3] i will add the right amount of a NaOH solution (KOH should be better) until it's fully dissolved than H2O2 again to get Cr(VI) as sodium chromate than finally crystallization.

Mildronate - 28-9-2012 at 11:00

My eyes burn. Resource spending, stainless steel is more valuable

blogfast25 - 28-9-2012 at 12:15

Salmo:

Your chloride ions are spectator ions: they don't take part in the reactions and are found in the solution part of the slurry.

Quote: Originally posted by Mildronate

My eyes burn. Resource spending, stainless steel is more valuable

Not if it's scrap it ain't. Salmo and others aren't trying to recover Cr from SS industrially, it really is just a bit of educational fun, as far as I'm concerned.

Endimion17 - 29-9-2012 at 03:58

Isn't chromium just an extremely tiny layer on stainless steel? Steel used for utensils and pots is usually iron-nickel alloy plated with chromium. It would take an awful lot of utensils to get small amounts of chromate salt...

I agree, if it's educational fun, go for it, but such chromate would be very expensive.

plante1999 - 29-9-2012 at 04:16

In fact S.S is S.S because of chromium, even cheapest S.S contain 18% chromium and iron without any other metal. Quality S.S contain 18% Chromium 15% nickel and iron. With dollar store spoon ( 18% chromium no nickel) it is not that costly. And with scrap it worth the cost to try my procedure.

Salmo - 29-9-2012 at 04:53

blogfast you say that Cl- ions are spectator ions, but if they are so, which are theyr plus charged counterparts?
In a solution this equation must be always true.
a[A+]+b[B+]+...= c[C-]+d[D-]+...
Anyway mildronate if you think that scrap stainless steel is more valuable than what i'm trying to do, you dont understand how much is possible to learn just with a spoon and some cheap OTC chemicals.

Endimion17 - 29-9-2012 at 05:04

Oh, so you mean not plated, but fully stainless steel? Because there are utensils that are plated only. They're cheaper.

Yeah, that might be very interesting, if we're talking about scrap.

blogfast25 - 29-9-2012 at 06:24

blogfast you say that Cl- ions are spectator ions, but if they are so, which are theyr plus charged counterparts?
In a solution this equation must be always true.
a[A+]+b[B+]+...= c[C-]+d[D-]+...

Ok. Let’s take the case of the oxidation and subsequent hydrolysis of FeCl2 to Fe(OH)3. Here are the balanced reactions:

1) Oxidation of the FeCl2 with peroxide:

FeCl2 === > Fe3+ + 2 Cl- + e-
½ x [H2O2 + 2 H+ + 2 e- === > 2 H2O]

Add up:

FeCl2 + ½ H2O2 + H+ === > Fe3+ + 2Cl- + H2O……. Eq.1

2) Subsequent hydrolysis of Fe3+:

Fe3+ + 3 H2O === > Fe(OH)3 + 3 H+……. Eq.2

3) Now add Eq.1 to Eq.2, scrapping redundant terms:

FeCl2 + ½ H2O2 + 2 H2O === > Fe(OH)3 + 2 HCl

(of course also: HCl + H2O === > H3O+ + Cl-)

Perfectly neutral, mon cher. Please note that above reaction equations represent stoichiometry (mass and electrical balances), not molecular reality! I could have written these reactions w/o even mentioning Cl and they would still have been correct: hence the term 'spectator ion'.

[Edited on 29-9-2012 by blogfast25]

Salmo - 29-9-2012 at 07:36

Damn you are right man I didnt thought about the "acidity" of FeCl3! I also found this on wikipedia:

Quote:

In industrial application, iron(III) chloride is used in sewage treatment and drinking water production.[13] In this application, FeCl3 in slightly basic water reacts with the hydroxide ion to form a floc of iron(III) hydroxide, or more precisely formulated as FeO(OH)-, that can remove suspended materials [Fe(H2O)6]3+ + 4 HO− → [Fe(H2O)2(HO)4]− + 4 H2O → [Fe(H2O)O(HO)2]− + 6 H2O

Anyway today i tryed to add NaOH to the slurry but it was impossible to understand what was happening in that solution..It was thick impossible to check the pH, i adedd more H2O2 and i tryed to filter what i could.. and it was impossible too believe me.
I will try again next week starting from other spoons and I will first precipitate everything (with NaOH) and than add H2O2 , for who wants to try too, remember never do the opposite

or a fucking brown colloidal shit will form.

[Edited on 29-9-2012 by Salmo]

blogfast25 - 30-9-2012 at 04:56

If it was colloidal it would probably run through the filter, like a solution. Fe(OH)3 has a tendency to do that and I've seen it happen: allof a sudden your precipitate ends up at the wrong side of the firlter, very annoying. But that didn't happen here... It's called peptisation. 'Sn(OH)4' does it too.

plante1999 - 30-9-2012 at 05:19

@ blogfast:

I think he would be better to use my procedure, more straight forward and more robust not to mention simpler.
He seam to have some difficulties with the peroxide method.

Salmo - 30-9-2012 at 08:29

Man i really dont know what it was, I just can tell you that it was a really thick sludge that wasn't forming any kind of precipitate, it was more like a fucking gel impossible to vacuum filter. Anyway as I'm bored to use nasty HCl at home

, tomorrow I will electrolyze a big spoon in a NaCl solution, I did it before with good results, then I will precipitate everything with NaOH, add peroxide and than filter the mass to obtain the chromate solution, i will post photos!
Plante your procedure is for sure better and faster I have to obtain some sodium nitrate first, can yuo tell me your exact procedure?

blogfast25 - 1-10-2012 at 09:46

@ blogfast:

I think he would be better to use my procedure, more straight forward and more robust not to mention simpler.
He seam to have some difficulties with the peroxide method.

To make such a statement a robust comparison between the methods would need to be made first. I started of with some kind of fusion (KOH + KClO3), only to find it wasn't necessary: everything can be done at RT and in solution/slurry.

An example of separating the Cr from an FeCr heating element, here:

http://www.sciencemadness.org/talk/viewthread.php?tid=21298#...

Still, why not post the details of your [plante] approach here?

[Edited on 1-10-2012 by blogfast25]

plante1999 - 2-10-2012 at 12:53

You have it right on the point that I should upload it, I would prefer to do a prepublication about it tough. I tried both and found the peroxide way a lot harder to do because of colloidals, concentrations of solution not to mention a lot more costly than hydroxide melt.

blogfast25 - 3-10-2012 at 07:30

Well, let us know when the prepublication is up.

tetrahedron - 12-10-2012 at 16:53

I also used electrolysis of S.S in a carbonate electrolite but I found it energy intensive and take time to make some chromate but Hasslefree.

sounds interesting..do you have any more details on this? do you add any oxidizer to the electrolyte?

use molten sodium hydroxide mixed with sodium nitrate for the chromium extraction from the dry mass

i just did the potassium version of this. after filtering out the largely unreacted Cr₂O₃ i got a yellow solution that i'll quantitate asap.

plante1999 - 12-10-2012 at 17:32

Only electrolysing S.S at 20V in Na2CO3 electrolyte. Sodium chromate go in to solution but iron and nickel don't. It take a large amount of time to get a product but it work really well. If molten for enough time all the Cr2O3 will react. Use stochiometry to know hydroxide/nitrate ratio.

12AX7 - 12-10-2012 at 19:46

If you have an oxidizer such as KNO3 or KClO3 on hand, you can perform a fusion:
http://webpages.charter.net/dawill/tmoranwms/Chem_Chromate.h...
The KCl and NaOH aren't necessary; with KClO3 at least, the reaction proceeds under "acidic" conditions, yielding KCl and K2Cr2O7 directly (but beware the fumes, exotherm and decomposition).

This may not work with ashed* stainless, because nickel or iron are probable catalysts for decomposition. I once tried the reaction on mixed oxides where chrome was known to be present; all I got was frothing (decomposition, O2) and a lot of black stuff.

It's noteworthy that molten KClO3 is reasonably stable in a steel or stainless crucible.

*By 'ashed', I mean oxidized, dried, powdered material. It could be burned in air (well, if you can about melt the stuff; it doesn't oxidize very fast otherwise), dissolved in acid and precipitated wholesale (mixed hydroxides), calcined, etc.

On a related topic, I seem to recall PbO2 can also be made by a similar method, but I'm not set up to test it.

Tim

blogfast25 - 13-10-2012 at 05:17

Hi Tim,

It's quite similar to treating Chromite (FeCr2O4) with KOH/KNO3, then leaching. The iron stays, the chromate is lixivated off. Nice page!

violet sin - 13-10-2012 at 05:58

so I have been messing around with the idea of purifying the chromite ore I found, but along the way just wanted to see if I could make Na2Cr2O7 from pure Cr metal. and I think I have come up with a pretty darn easy way.

I made a decently strong NaOH solution and immersed a ~99.1% chunk of ebay Cr metal(+) electrode and a Ti(-). the solution turns bright yellow after a short while, because of deep orange gel pouring off the Cr electrode. this worked so far... the colors were right. I figured ok let it run a while build up a concentration. it developed a darker(greyish) cast and became more opaque. My only guess was that there was no longer any sodium to keep up with the dissolving Cr. there were some small fragments of Cr metal so I figured maybe I can add H2O2 and or lye to see if I could oxidize it after providing more Na ions. both caused bubbles and heat. after filtering w/ coffee filter I found a grey gel with the tiny bits of Cr glimmering in it, and a BRIGHT yellow clear solution. this was good for the chromate I thought, and after adding a small amount of H2SO4 drop-wise to the solution it became decidedly orange! again I figured I was on the right path. the H2SO4 didn't noticeably fight with NaOH as no bubbles were evolved. figured this meant the sodium was used up.
I put this flask in a water bath just below boiling, and waited patiently. went from 125ml --> ~50 ml before I noticed anything. orange crystals at the edge of the liquid. filtered it around 30ml and got a decent amount of orange crystals. though they don't seem to be too regular. need to recrystalize after washing w/ ice water. my 12th merck said that sodium dichromate aq. can be set to ~100 for a prolonged time to make anhydrous. I deff had crystals and didn't see any breakdown, and the onely elements I introduced were Cr, NaOH, H2O, with a bit of H2SO4, and H2O2. so given the behavior and colors I figure I have Na2Cr2O7

though tired now and ready for bed 6:30am...

but seems pretty simple to use NaOH sol. and slap a Cr electrode in there. then throw a small amount of juice at it. my cell(250ml Erlenmeyer) needed to be put in a buffering water bath. not cooling but to tame against thermal shock. my power supply was an OLD hacked ATX computer spare.

maybe tomorrow I can figure out how to get some pic's up. gotta buy some more gloves also. if some one thinks I may have performed some other reaction then please enlighten me. but I feel fairly confident. if it did work I don't see why this couldn't be adapted to the use of Stainless Steel. though I would stay the heck away from HCl with this stuff. chromyl chloride isn't for me.... I know that in the test run of this I used a stainless nail to make contact to the Cr metal. the Fe contamination just turned to a brown gelly at the bottom, and man does Ni drop out of solution with NaOH in there. ( just got done making nickel chloride, sulfate, acetate and hydroxide last week). any thoughts?

[Edited on 13-10-2012 by violet sin]

Salmo - 13-10-2012 at 06:43

Ok finally I did it!
I told you that I would have posted my procedure and the photos so here I am.
I started from 18/10 stainless steel from a fork, I electrolyzed it in a 25% NaCl solution with my lab power supply adding some 33% HCl to destroy the hydroxides formed. than I stopped it after 2 hours and I filtered the solution to obtain a clear green solution.
Than I added a saturated Na2CO3 solution to precipitate iron/nickel and chromium as hydroxides, I didn't use a NaOH solution because strong sodium hydroxide solutions react with chromium hydroxide to form soluble [Cr(OH)6]3-.
Anyway after the precipitation I filtered the solution and I obtained a dark green/brown sludge that I washed with water. Than heating the sludge, I added a really strong NaOH solution to it,and than 35% H2O2, 10ml a time, to oxidize Na3[Cr(OH)6] to sodium chromate, so the sludge after lot of fizzing, turned brown and I filtered it obtaining a yellow sodium chromate solution.
So because I wanted to remove the excess of NaOH, NaCl and everything else I added HCl to the sodium chromate solution until acid, I reduced it adding hydrogen peroxide again and I precipitated Cr(OH)3 adding sodium carbonate.
I washed my chromium hydroxide with hot water many times and I obtained a blue mass that I redissolved with as little as I could of a strong NaOH solution than I added the right amount of H2O2 and I heated the solution, after filtering I got a really strong yellow sodium chromate solution.
Now I checked the pH of the solution and it is high so I have some sodium hydroxide in it, do you know how can I remove it for the last time?? I thought to sulfuric acid until pH 7 maybe sodium sulfate is easy to remove.. I dont know
Damn I want an induction heater and damn believe me filtering hydroxides or oxides is time consuming and a pain in the ass.

[Edited on 13-10-2012 by Salmo]

blogfast25 - 13-10-2012 at 12:14

Well done, Salmo.

Two comments:

Than I added a saturated Na2CO3 solution to precipitate iron/nickel and chromium as hydroxides, I didn't use a NaOH solution because strong sodium hydroxide solutions react with chromium hydroxide to form soluble [Cr(OH)6]3-.

[snip]

Now I checked the pH of the solution and it is high so I have some sodium hydroxide in it, do you know how can I remove it for the last time?? I thought to sulfuric acid until pH 7 maybe sodium sulfate is easy to remove.. I dont know

Your reasoning is correct but chromite [Cr(OH)4(-)] does not appear to form when large amounts of iron are also present (as is the case here). In those circumstances, the Cr3+ appears to co-precipitate fully with the iron oxide(s). But as you’ve seen later on, Cr(OH)3 on its own does dissolve in strong NaOH to form sodium chromite.

Re. removing the remaining peroxide. DO NOT acidify at this point because in acidic conditions the excess peroxide will reduce the chromate back to Cr (III)! The best way of destroying left over hydrogen peroxide is to simmer your solution for some time. Then take attest tube sample and acidify it. It should revert to the yellow/amber of dichromate. If any blue/green appears, you’ve still got peroxide left.

Re. removing the excess NaOH. For dichromates usually the K salt is preferred, I believe because K2Cr2O7 is only sparingly soluble in cold water and thus easier to isolate. The excess KOH would be neutralised with H2SO4 and on cooling crystals of K2Cr2O7 would appear, provided you started from about the right amounts of KOH and Cr. I’m not sure how this will pan out with Na2Cr2O7.

Nice photos!

[Edited on 13-10-2012 by blogfast25]

violet sin - 13-10-2012 at 22:17

so was there anything wrong with the route I took? I don't have Ph test strips and reagents are quite limited.

composit.bmp - 728kB

the crystals aren't defined well in the cigar tube b/c I didn't have time to do anything but give 'em a quick rinse in the filter. were there any concerns of contamination? other than Na2SO4? blogfast25 you said something about soluble chrome hydroxides but not to worry in an alkaline setting? just been having fun with the electrochem and a few of the metals I collect.

blogfast25 - 14-10-2012 at 05:13

VS:

Is that sodium dichromate or potassium dichromate in the cigar tube? To purify further, recrystallise.

Nice work!

Re. Cr(OH)3 and chromites, see further upthread...

Salmo - 14-10-2012 at 06:08

Thanks blogfast today I will acidify drop by drop with sulfuric acid to ph 7 and than I will try to separate the Na2SO4 using its different solubility.
Na2SO4 is 4.9g/100ml ( at 0°C) and sodium chromate is 163g/100ml ( at 0°C).
Violet sin, I think you have my same problem, the fucking sodium! anyway good job now recrystallise.
Check this site: http://www.chemguide.co.uk/inorganic/transition/chromium.html

[Edited on 14-10-2012 by Salmo]

violet sin - 14-10-2012 at 08:57

blogfast: thanks that is the Na salt. obviously not pure as I am reading in "atboic:Cr" from the reference section right now. thought the stuff in the cigar tube is from the first stuff to fall out from water bath reduction, and more likely pure(er)...

facing Cr(OH)3 and the chromites it makes with NaOH... gotcha

from what I gather you are saying, starting off with a strong NaOH sol for electrolyzing means I have the Na3[Cr(OH)6] and Na[Cr(OH)4] from the moment Cr oxidizes into solution? -but- it can be oxidized from the -ite to the chromate with H2O2 in it's alkaline situation. then further boiling I would get rid of extra hydrogen peroxide, before H2SO4 acidification to the dichromate( which I wanted for conversion ). but I only really have a bit of Na sulfate salt contaminant from neutralization of excess NaOH *IF* I managed to get all of the odd hydroxide complex completely oxidized to chromate first *AND* I boiled it long enough to kill the H2O2 before acid so it doesn't revert to the chromite correct?
the grey gel I had with the sparkly bits in it may be the hydroxide 'aging out', and as a result of Cr3+ from adding the H2O2 in acidic conditions? explains why it got cloudy a few times( playing with H2O2)

or is there still a chromite that is not solved by oxidation to chromate with H2O2? I only have 3% so it is a LOT of fun to add and boil down. if there was any left over then it is surely gone by the time I through with it's water... those older books are kinda hard to get the hang of,, having to rearrange stuff in your head for hydrated forms or odd oxides that are a mix of two other oxides. I say if you haven't been doing that for a while it's a real treat to try and keep it all straight

thanks also salmo nice info site. like it. I think the reason for doing it in Na form is the merk says it will be fine to dehydrate --> anhydrous by water bath or @ ~100'C for extended time w/o decomp probs. then it would be easy to add water necessary and do the KCl dance quickly. that's my end goal with this exploration, no reason just looks fun. I just wish I could accidify it with something that didn't have a fun salt leftover like Na2SO4

[Edited on 14-10-2012 by violet sin]

12AX7 - 14-10-2012 at 14:09

Chromates are very strong in color and stain everything, so you could have sodium sulfate with surprisingly little (di)chromate. Not sure if they form solid solutions; I suppose chromate (but not di-) could partially substitute sulfate.

One easy way to tell, take some crystals and attempt to dry or melt them. Sodium sulfate hydrate decomposes under low humidity conditions, or melts in its water of crystallization, IIRC around 80C.

Precipitating with potassium has the downside that K2SO4 is also fairly low solubility. Ditto ammonium.

Using HCl instead of H2SO4 would be better; both KCl and NaCl have good solubility at all temperatures.

Tim

violet sin - 14-10-2012 at 15:50

well it would be hard to have too much Na2SO4 in there b/c I only had to use a few ml of H2SO4 to acidify to orange, whereas my chromium metal chunk went from a nice angular shattered look to a well worn beach pebble. I think in the entire 250ml flask I used ~5ml for shift(<10ml, was only a 1ml eye dropper). that leads me to believe that I used up most of the NaOH forming the sodium chromate and chromite that got oxidized so ya chromate( as I used 4-5 Tbs. NaOH). I stopped energizing once I noticed a greyish opacity(Cr(OH)3), then filtered. that was more than likely the solution ppt out Cr(OH)3 because of lack of free OH- to solovate. then oxidized and acidified. so I was deff on the right track but had the hydroxide/chromite like blogfast n salmo said. not saying you are wrong just I don't think I had a lot of sulfate...

looking back after the last few very helpful posts, I dissolved everything again this morning and added H2O2 back to clear yellow. only a small pile of clean Cr crumbs were left in the eddy after much bubbling and quite a few additions. seemed I had a rather sizable amount of Cr(OH)3/chromite. I left it to simmer @~90'C for like an hour to boil down and kill H2O2.

I did notice when dissolving crystals again that something was wrong, took more water & time than it should have, so quite glad I cleaned it up. I like the dry melt idea, shoulda looked up more physical traits to compare before declaring DONE. good point. no HCl for me with Cr, afraid of the chromyl chloride. I have less than ideal ventilation... no thanks. working on a dimmer-switched vent fan/hood and ducting today as well. construction background comes in handy yet again

trying to get some pictures up for re-polished product soon. also started a small batch from stainless nails. doing everything salmo did just on a 50ml scale just for shitts and giggs. salt works surprisingly well for devouring SS

12AX7 - 14-10-2012 at 17:11

Yup, if there's a lot more chrome than sulfate then it's probably not going to show up.

Chromyl chloride isn't a concern, as far as I know: it hydrolyzes rapidly in water. A bigger concern would be dichromate oxidizing chloride to chlorine and reducing your yield, but I think this is unfavorable too, and in any case not a problem if an excess of acid is not added.

Tim

Salmo - 15-10-2012 at 04:24

Today I will try to purificate my sodium dichromate solution removing all the sodium sulfate inside, i will boil down the solution to 10ml than I will put it in the fridge to 0°C hoping to see some sodium sulfate precipitation but anyway it's hard that i will obtain something pure..i think that solid fusion is the best idea because it's so difficult to add the right amount of NaOH during the oxidation and adding too much means a not pure product full of NaCl or Na2SO4..

[Edited on 15-10-2012 by Salmo]

blogfast25 - 15-10-2012 at 09:24

Yes, separating the chromate/dichromate from neutralisation by-products may prove to be the hardest part here.

Obviously, however you want to make these Cr (VI) salts, the trick is to keep by-products to a minimum by working as close to stoichiometry as possible. Small amounts of soluble by-products should then always remain in solution. So as with any decent synthesis, planning ahead is crucial.

Tim’s right about the solubilities, KCl is somewhat more soluble than K2SO4. HCl also has the advantage that it’s volatile so that any excess will evaporate on drying your end-product.

The solubilities for the Na and K dichromates at 0 C (Wiki solubility table) are resp. 163 [that sounds really high!] and 4.7 g/100 ml, so for dichromates the alkali of choice should be KOH,H2O2 as oxidiser and HCl as neutraliser. Unless you want to use fusion.

[Edited on 15-10-2012 by blogfast25]

tetrahedron - 15-10-2012 at 09:50

Only electrolysing S.S at 20V in Na2CO3 electrolyte. Sodium chromate go in to solution but iron and nickel don't. It take a large amount of time to get a product but it work really well.

i've been running this the whole day, using a chrome plated stainless steel fork for anode and a graphite cathode. the chrome plating resisted corrosion so well that i had to temporarily hike the voltage to almost 40V before the electrolyte caught color (a previous addition of H₂O₂ didn't help). i'm finding it hard to keep the voltage that high as my PSU maxes out at 6A. i had to increase resistance by reducing the anode surface dipping in the electrolyte, but this feels wrong from an efficiency standpoint. does it necessarily have to be that high to oxidize the Cr to +6?

12AX7 - 15-10-2012 at 16:16

Hmm, K2CrO4 has higher solubility than K2Cr2O7. Might be useful. pH swing crystallization or something? Take your impure mixture, add an estimated stoich amount of KOH, wait to dissolve; evaporate to saturation, then add HCl. KCl and NaCl stay put (unless they are also at saturation, in which case the added chloride may drop a small amount; a second pass should be clean).

And, of course, collect all wastes and crystallize out all the orange and not-orange stuff and repeat...

Tim

plante1999 - 15-10-2012 at 17:19

Only electrolysing S.S at 20V in Na2CO3 electrolyte. Sodium chromate go in to solution but iron and nickel don't. It take a large amount of time to get a product but it work really well.

Like I already said, It take time is inefficient , but It work. If your S.S is hard to dissolve use Sodium bicarbonate electrolyte, It should help for the dissolution of the Chromium. The idea is that the Ph is lowered increasing corrosion of iron.

tetrahedron - 16-10-2012 at 03:45

after almost 10h operation, the fork was only superficially corroded (at the teeth). meanwhile, a stratified sediment formed, which was removed after cooling:

the tip is undissolved soda. i think the layers are due to fluctuations in throughput due to the solution shrinking below the anode level, and later being filled up.

after the convex side dried up, the mass was turned over. the flat side was covered in brown/purple crystals, like a geode. unfortunately i wasn't able to capture the crystalline look, as the residual soda solution quickly dried and covered it in white too:

sadly this was the most interesting part of the experiment. the leftover electrolyte has almost completely solidified, the filtrate looks a rusty brown. i'll try to verify if there's any chromate in there.

edit. no appreciable change with thiosulfate.

[Edited on 16-10-2012 by tetrahedron]

blogfast25 - 16-10-2012 at 07:41

Quote: Originally posted by 12AX7

Hmm, K2CrO4 has higher solubility than K2Cr2O7. Might be useful. pH swing crystallization or something? Take your impure mixture, add an estimated stoich amount of KOH, wait to dissolve; evaporate to saturation, then add HCl. KCl and NaCl stay put (unless they are also at saturation, in which case the added chloride may drop a small amount; a second pass should be clean).

And, of course, collect all wastes and crystallize out all the orange and not-orange stuff and repeat...

Tim

That's an interesting thought, Tim. For it to work, good control of concentration would be required. But it would be great to see the dichromate drop out on acidification!

Salmo - 16-10-2012 at 11:49

Ok my experiment failed today on purification, but as a good scientist

I didn't give up and today I'm starting again.
Maybe chromates are not so interesting compounds because there are better ones but I want to develope the right procedure because I love OTC chemistry and I think that chromates could be the cheap choice for the poor home chemist (like me) to do his oxidations.

Anyway blogfast you're right as always, stoichiometric calculations and quantities are absoluteley necessary with this procedure unfortunately I understood this just today!
Today I started the dissolution of a 40g 18/10 stainless steel in 220ml (I'm using an excess) 37-38% H2SO4 (5M)

These are the calculations:

Iron: 28.8g (0.516mol)
Nickel: 4g (0.068mol)
Chromium: 7.2g (0.138mol)
SS dissolution in H2SO4 37-38%:
Fe + H2SO4 -> FeSO4 + H2
Ni + H2SO4 -> NiSO4 + H2
2Cr + 3H2SO4 -> Cr2(SO4)3 + 3H2
TOTAL H2SO4 moles necessary: 0.791 TOTAL H2SO4 37-38% mL necessary: 160ml

Then before precipitation I will first boil down and than filter the dark green solution to remove unreacted stuff like carbon.

Precipitation with NaOH:
FeSO4 + 2NaOH -> Fe(OH)2 + Na2SO4
NiSO4 + 2NaOH -> Ni(OH)2 + Na2SO4
Cr2(SO4) + 6NaOH -> 2Cr(OH)3 + 3Na2SO4
TOTAL NaOH moles necessary: 1.582 moles TOTAL NaOH grams necessary: 63.28g

Then I will filter and wash many times the green precipitate to eliminate 102.55g of Na2SO4, so I calculated that I will need 16.56g (of NaOH to make this reaction happen:
[Cr(H2O)3(OH)3] + 3OH- -> [Cr(OH)6]3- + 3H2O
(I hope that the iron hydroxide precipitate won't hold too strongly the chromium hydroxide anyway I will heat the sludge for some time).

So I will add an unknown amount of 35% H2O2, it's hard to calculate because of hydrogen peroxide decomposition and Iron hydroxide oxidation reaction that takes place at the same time, the Schikorr reaction:
3 Fe(OH)2 -> Fe3O4 + H2 + 2 H2O

Maybe using less than 16.56g of NaOH could be better I mean it would be less risky even if the yeld would be lower, I think I will use 15.00g instead. But wait, there is another problem.

This is the reaction for the oxidation of Cr(III) to Cr(VI)

2[Cr(OH)6]3- + 3H2O2 -> 2[CrO4]2- + 2OH- + 8H2O

As you can see the oxidation reaction produces 2 moles of NaOH for each mole of chromate produced and this is not good because we will have a lot of salt to remove in the purification stage and this is really hard.

So we have to think about it, maybe the key is to lower the NaOH quantity because some would be generated “in situ” from the oxidation but is that possible to calculate? I don’t know I just know that I have to put 3 moles of sodium hydroxide to produce 1 mole of sodium chromate and 1 mole of sodium hydroxide. Maybe we have to use 2/3 of the initial quantity? So 11.04 grams? I would put 10 grams to stay safe and I would do the dissolution/oxidation while heating and stirring well the sludge/solution because of the lower initial concentration of sodium hydroxide.

What do you think?

[Edited on 16-10-2012 by Salmo]

tetrahedron - 16-10-2012 at 17:56

If molten for enough time all the Cr2O3 will react. Use stochiometry to know hydroxide/nitrate ratio.

what's the ratio in your "drano"? i used stoichiometric amounts according to the reaction

Cr₂O₃ + 3KNO₃ + 4KOH ---> 2K₂CrO₄ + 3KNO₂ + 2H₂O

but i read elsewhere that the nitrate acts as a catalyst for air oxidation, so the nitrate can actually be less, right?

i repeated the experiment with a dilute Na₂CO₃ electrolyte this time. Again I can't say for sure it was successful, although the higher voltage did help in corroding the stainless steel somewhat. OTOH in a diluted NaCl electrolyte the anode corroded quickly and efficiently, but no chromate was observed.

violet sin - 16-10-2012 at 19:08

salmo: ok so what you are saying is... the 3Na+ in solution gets absorbed with the [Cr(OH)6]3- ion. but upon oxidation only 2 are taken up by the chromate and/or dichromate. leaving one mole in solution. to cancel. which leaves a 1:1 ratio of leftover Na2SO4 per each Na2Cr2O7(or Na2CrO4). means a 50% contamination. LAME!!

but check this from melcor in ref section:
J. Fritzsche added a warm soln. of potassium
dichromate to an excess of conc, sulphuric acid ; the chromic anhydride separates
in small red crystals. The liquid is drained from the crystals, which are then
dried on a porous tile over sulphuric acid. The crystals are then recrystallized
from an aq. soln. P. A. Bolley said that the chromic anhydride so prepared
contains a little sulphuric acid as impurity and the Metals Protection Corporation
removed the sulphate by means of barium hydroxide, carbonate, or chromate.
R. Bunsen, A. V. Rakowsky, A. Dalzell, F. Dietze, 0. Ficinus, H. Moissan

so it seems that if you get a decent dichromate sol (even if it has Na2SO4 from acidification i think). treat sol w/ some con. H2SO4 and ppt. red chromic anhydride crystals. (weird cause CrO3 is solubile in H2SO4?) with still totally soluble Na2SO4?.... effectively separating with minimal/acceptable losses. resolovate with min H2O then drop out remaining H2SO4 contamination w/ barium as BaSO4 ppt. but don't overshoot. filter you aq. chromic acid solution free of Ba2(SO4)3 (aq sol, 0.00024g/100ml @25'C) and dehydrate or not... with pure chromic anhydride, make chromic acid all you would have to do is add a carbonate of K/Na and let the CO2 gas off right? sounds like the button to win, so I'm gonna press it. have a lill bit of NO3- left in sol but decomp heating should kill it

I am going to try a 5-10ml sample tonight. going to be last experiment for a while

Salmo - 17-10-2012 at 02:40

You had a nice idea violet sin but sadly I've no con. Sulfuric acid.. And another thing I hope that no Na2SO4 would co-precipitate I mean I don't know its solubility in con.sulfuric acid.. Anyway why not to try?
Today I will try to use math to solve the NaOH problem even if I think that using 10g or less like 8 would solve the problem without loosing precious time

what I think is that there is a consumption of sodium hydroxide so the reaction stops at a certain time and all the NaOH is over.

violet sin - 17-10-2012 at 05:55

after putting a few ml of con H2SO4 in a testtube, I added a few drops of sodium dichromate. my straw color ACE hardware acid turned a smokey red, and yet clear top layer. 1/8" thick. but when I cam back to it a while later I found a bright red opaque line on the side of the tube at solution line. it was bright red and kinda looked like tiny particles of sand. got a pic but not time to put it up. it was a very verry small amount. but clearly visible. and unlike any other sulfate or chromate/dichromate I have seen to this point. so I believe it was indeed chromic anhydride. I tried to fish some out to try a drop of MeOH on it(saw youtube vid w/ CrO3-EtOH = flames). I could feel it crunching as I tried to get it out. but most stuck to the testube wall.

it seems to be produced and if shaken dissolves again. I noticed this bright opaque lipstick colored ppt on initial addition of the dichromate sol. but shook it and turned into a red/rootbeer colored band. may be tiny dispersed crystals. I retried and got the above results. with the red line by not disturbing it.

also from one of my last attempts to ppt the dichromate pissed me off, so I returned it to chromate boiled down(in prep to oxidize again) a bit and left stand for a day while trying to figure it out. came back to BIG LONG bright yellow crystals. sever were 3" long and roughly 2/3 the diameter of a standard pencil. I broke one open see if it was just sodium sulfate coated with a vibrant chromate... but no it was yellow all the way through. I know I had a lot of Na2SO4 in there, but it seems to have formed another complex with the chrome ions. again I have pictures but will have to wait till later today to retrieve them from my phone resize and post.

doesn't seem like it would be feasible to do for small scale with a decent return, and with large scale you have a LOT of hexavalent Cr and con. H2SO4 laying around... but was worth a try. would like to see some large CrO3 some day.

DerAlte - 17-10-2012 at 09:43

There seems to be an increased interest in the forum recently re Chromium compounds. I have collated the segments I wrote in this thread into a single document and made it available as

http://www.sciencemadness.org/scipics/Chromium%20Revisited.d...

I have not changed the content except for a few typos. Any original errors will still exist (there are always some!). Don’t go copying it for term papers, etc. It may not be academically correct. For amateur use only! Always read several references, not just one…

Regards, Der Alte

tetrahedron - 17-10-2012 at 12:08

thank you, a lot of great info

blogfast25 - 17-10-2012 at 12:29

Salmo:

Forget about the Schikorr reaction, it doesn’t happen in those circumstances. But it helps to get the stoichiometries right (here with K instead of Na):

For the oxidation of the chromic hydroxide to chromate:

Cr(OH)3 + 2 KOH + 3/2 H2O2 === > K2CrO4 + 4 H2O

For the conversion of the chromate to dichromate with HCl:

K2CrO4 + HCl === > ½ K2Cr2O7 + KCl + ½ H2O

Overall:

Cr(OH)3 + 2 KOH + HCl + 3/2 H2O2 === > ½ K2Cr2O7 + KCl + 9/2 H2O

So bearing in mind that you’re also alkali oxidising Fe2+ to Fe(OH)3 , estimating how much H2O2 you need isn’t difficult. Use a calculated excess (e.g. 20 % extra) and add to chilled slurry to avoid losing too much H2O2 to simple decomposition. Then simmer the slurry a bit to decompose residual peroxide. Filter to extract chromate.

Interestingly it also shows that if dichromate is the goal, an accompanying K or Na salt is always inevitable.

[Edited on 17-10-2012 by blogfast25]

Salmo - 17-10-2012 at 12:46

I'm an idiot

I didn't see that you (blogfast) merged the two equations in one anyway I think that you are right it is simpler than i thought about the sodium hydroxide, in one two days I will finish my fourth experiment about chromates.. today I finished the dissolution of my 40g fork in sulfuric acid I will post a useless photo in some minutes!

Thank you really much DerAlte really useful document.

[Edited on 17-10-2012 by Salmo]

[Edited on 17-10-2012 by Salmo]

tetrahedron - 18-10-2012 at 11:52

yesterday i took some of the sludge that i got when "electrodissolving" stainless steel in a Cl^- containing electrolyte and poured it into a fresh solution of NaOH and NaCl (about 1 spoonful each in 400ml tap water), which i then electrolyzed for a couple hours with a 5mm diameter graphite "gouging rod" anode, initially @ ~7V, soon reduced to 4.4V & 3A in order to reduce the loss of chlorine. i noticed a faint chlorine smell during the procedure (it was done outside), as well as after unplugging the current. even after standing overnight i still noticed some bubbling, which i can't explain.

the chloride only acts as a catalyst, first converting at the anode into chlorine and immediately afterwards to hypochlorite, which oxidizes Cr(III) to Cr(VI), reverting to chloride.

the filtrate looks promising

although i still don't have a way to quantitate the chromate present, and i can't tell how much chromite is contributing to the color.

[Edited on 18-10-2012 by tetrahedron]

blogfast25 - 18-10-2012 at 12:43

Tetrahedron:

Have you tested for chrmate with a lead (II) salt solution?

Salmo - 18-10-2012 at 12:43

What do you mean when you say chromite?
anyway your method is nice even if you will have many impurities. You first generate chromium chloride than you precipitate it and than some is dissolved from some hydroxide than the "in-situ" generated sodium hypochlorite oxidize Cr(III) to Cr(VI).. ok but you know that a yellow color doesen't mean anything about the quantity of sodium chromate present? Anyway add some acid and look if your solution turns orange.

[Edited on 18-10-2012 by Salmo]

tetrahedron - 18-10-2012 at 13:34

Have you tested for chrmate with a lead (II) salt solution?

no, unfortunately i don't have any lead salts. i guess calcium could also work, but it'll have to wait.

What do you mean when you say chromite?
anyway your method is nice even if you will have many impurities. You first generate chromium chloride than you precipitate it and than some is dissolved from some hydroxide than the "in-situ" generated sodium hypochlorite oxidize Cr(III) to Cr(VI).. ok but you know that a yellow color doesen't mean anything about the quantity of sodium chromate present? Anyway add some acid and look if your solution turns orange.

thanks, however allow me a correction. i don't generate chromium chloride (because the first electrolyte is not acidic), but chromium oxide/hydroxide (along with the corresponding Fe and Ni compounds, and carbon, and who knows what other additives are in steel cutlery), which precipitates out forming the aforementioned sludge.

by "chromite" i mean the soluble anion CrO₂^- given by the Cr(III) in alkaline solution. i don't know what this looks like, probably lime green.

of course the color can be deceiving..however i compared it to the leachate of the alkaline melt method and i made a 50% educated, 50% wishful guess that i got it right, and called it a day ;p

[Edited on 18-10-2012 by tetrahedron]

elementcollector1 - 18-10-2012 at 15:21

I've tried reducing my Cr(III) solution (reduced with acid+alcohol) with baking soda + boiling water with no results. I think I'm going to suck it up and try boiling the emerald green solution down, then adding weak NaOH to get my Cr(OH)3.

tetrahedron - 18-10-2012 at 15:31

I've tried reducing my Cr(III) solution (reduced with acid+alcohol) with baking soda + boiling water with no results. I think I'm going to suck it up and try boiling the emerald green solution down, then adding weak NaOH to get my Cr(OH)3.

no no no..oxidize! you're in the wrong thread ;p

elementcollector1 - 18-10-2012 at 16:33

no no no..oxidize! you're in the wrong thread ;p

...Dangit!

violet sin - 18-10-2012 at 21:20

hey here are the pictures I meant to post

red chromic anhydride from the warm dichromate dropped into concentrated H2SO4. the tube was laid to rest for a while tipped a bit and made the opaque crystals.(lill red upside down smile)

here are the large crystals that grew over a day or so when I got mad about the Na sulfate contamination. not too bad for growing that quick in a 200ml flask. still not 100% sure what the composition is though.

here is most the pile in a watch glass next to a dime for size comparison.

sorry the pic's aren't the best only off my phone. but *seems* that the chromic anhydride description from the reference book worked just fine.

Salmo - 18-10-2012 at 23:10

Hey violet wonderful crystals ! I think they could be pure

blogfast25 - 20-10-2012 at 06:32

Quote: Originally posted by violet sin

Wonderful crystals indeed, but what are they? The solubilities at 20 C (Wiki solubility table) of sodium chromate and sodium sulphate are respectively 84 and 19.5 g/100 ml. If sodium sulphate is present in appreciable quantities it could well crystallise out first.

You really need to take some of your crystals, crush them up and dissolve them in a minimum of hot water, then allowed to crystallise again and see what they look like…

blogfast25 - 20-10-2012 at 10:32

Meanwhile I’ve done another chromate/dichromate experiment, this one aimed at staying close to stoichoimetry outlined above and keeping the solution as concentrated as possible, with a view of possible precipitating the Cr as dichromate in acidification.

1/10 mol of ‘Cr(OH)3’ was prepared from chromic sulphate hydrate by precipitation with NaHCO3, filtering and washing the green precipitate. The obtained paste, about 167 g (containing about 10.3 g of ‘Cr(OH)3’) was quantitatively transferred into a 500 ml beaker.

To it was then added 14 g KOH (about 10 % excess and accounting for water in the KOH) dissolved in 15 ml water. The paste partly dissolved and a thin, green-blue slurry of potassium chromite/’Cr(OH)3’ was obtained.

To this slurry, chilled 35 % H2O2 was added from a 10 ml burette, very slowly. Stoichiometrically about 15 ml would be required. At first everything went well: the slurry gradually changed colour (from green blue to khaki brown), it also thinned further and the colour of the liquid phase (on my glass stirring rod) was yellow. There was modest effervescence and gradual warming of the slurry. But after adding 30 ml of 35 % H2O2 the slurry still hadn’t cleared up fully (there should be more than enough water to carry all the formed potassium chromate) and I stopped adding peroxide. Final volume was by then about 200 ml, temperature about 45 C and pH about 7 – 8 (by paper).

I then simmered the slurry for a bit and allowed it then to stand and cool. Below is a photo of it after about 15 minutes of standing:

The supernatant liquid clearly shows there is chromate in solution (although at that pH perhaps chromate/dichromate in equilibrium) but there is also a considerable amount of an unexplained brown, heavy (it sinks pretty fast) precipitate. Note that bad lighting shows the liquid far darker than it is to the naked eye. It's really strong yellow.

The fairly low pH at the end is of course a bit worrying: did oxidation stop because of lack of alkalinity? How then to explain the brown precipitate?

Tomorrow the supernatant liquid will be decanted off and some test tube tests will be done on the precipitate.

[Edited on 21-10-2012 by blogfast25]

Arthur Dent - 20-10-2012 at 10:45

Yup, as blogfast25 mentioned, dissolution and recrystallisation of those lovely crystals will probably help you isolate a product that's more pure, as well as give you a good idea of the solubility of the crystals you grew, so that you could end up with a cleaner product after a few recrystallisations . Sodium sulfate crystals should be snow white, and sodium chromate crystals are a bright, intense canary yellow, and should form before the sulfate crystals.

BTW glad to see you're back 12AX7, it's been a hell of a while!

Robert

elementcollector1 - 20-10-2012 at 11:28

Sodium bicarbonate works too?
Anyway, I had some Cr(III) in solution and a small chunk of NaOH was tossed in; it all effortlessly turned yellow. No H2O2 was needed...

blogfast25 - 20-10-2012 at 11:51

Sodium bicarbonate works too?
Anyway, I had some Cr(III) in solution and a small chunk of NaOH was tossed in; it all effortlessly turned yellow. No H2O2 was needed...

Whatever happened, it can't be oxidation of Cr3+ to chromate: that really does require a powerful oxidiser. Tst for chromate with a lead (II) salt (yellow PbCrO4 precipitates)

Yes, in many cases NaHCO3 is a strong enough alkali to precipitate a hydroxide. In some case you'll obtain a basic carbonate though. Cr doesn't appear to form one of those though...

violet sin - 20-10-2012 at 13:33

as I said before I broke open the crystals to make sure they were not clear inside. and they were not. just a beautiful citrine(gem) color. but I did notice (and forgot to mention)that when left in air for a while(15min or >

they turned a bit frosty white on edges. reminded me of HCl frosting stuff around when evaping CuCl2 or the like( didn't have any acid out or heated salts though).
also I am ~4 hours away from home for the next 2 months working so I can't answer any of your inquiries on the crystals until I can set foot in my lab(aka the basement) again

i will do some investigation for sure and post my findings.

oh ya, another thing I noticed... when I filtered the solution off, it was still quite yellow but noticeably less so than before. like it might have taken a decent quantity of chromate with it.

one can only hope it is all chromate

also, seeing as how it was filtered while hot to remove flocky light colored stuff after H2SO4 addition. then left to sit n crystalize, after which I relieved the solution of large yellow crystals, I got more of it( flocky ppt). thinking concentration dependent release by say common ion Na+1 but sol was not green or violet so chromate wasn't reduced or ligand bound by SO4-2 *to the best of my knowledge*

[Edited on 21-10-2012 by violet sin]

Salmo - 21-10-2012 at 03:00

I think that if you want to be sure you should make a test with some Barium chloride or calcium chloride, if no precipitate it should be pure enough.

blogfast25 - 21-10-2012 at 05:23

I think that if you want to be sure you should make a test with some Barium chloride or calcium chloride, if no precipitate it should be pure enough.

The test for sulphate ions with a Ba salt solution would be best carried out on the recrystallised product because the test is so sensitive.

blogfast25 - 21-10-2012 at 08:33

So here are a few experiments I conducted with the products obtained above.

The supernatant liquid was carefully decanted off and the brown insoluble residue caught on a filter and washed free of (presumed) potassium chromate solution.

To the presumed potassium chromate solution (150 ml) was added 10 ml of HCl 36 %. Colour firstly darkened a lot, a bit of fizzing occurred and then the colour settled after a few seconds to the familiar amber of dichromate. I estimate the solution to be about 0.1 M in K2Cr2O7. About 1 ml of solution was further acidified with H2SO4, then MeOH was added and on steam bath the green colour of Cr (III) appeared quickly, confirming the amber colour is caused by K2Cr2O7 (oxidation of a primary alcohol to its coresponding carboxylic acid).

The insoluble residue had in the mean time acquired a slightly greenish hue, less dark than yesterday. A good pinch of the filter cake dissolved very quickly in hot 36 w% HCl to green Cr(III). With hot KOH (unknown but strong) dissolution of a similarly sized pinch to green chromite (III) took a little longer but happened too. Both tests were done in test tubes. At NO POINT did I observe gas evolution (oxygen or chlorine) which if observed would have been indicative of either a peroxo complex or an unexpected (higher) oxidation state of the Cr.

The chromite test tube was then cooled under running tap water and an arbitrary amount (a few ml) of 35 % H2O2 was slowly added in two lots (cooling in between) and the oxidation to chromate took place swiftly.

Conclusion: in the absence of other data I must conclude that the failure to oxidise and dissolve all ‘Cr(OH)3’ was due to lack of alkaline reserve and that the unusually coloured residue was in fact nothing else than ‘Cr(OH)3’ (or a hydrated form of chromic oxide).

I might have another go at this, this time using 3 mol of KOH per mol of ‘Cr(OH)3’ instead of the previously used 2 mol/1 mol ratio.

Had the synthesis proceeded as planned (and at 100 % actual yield), a solution of about 7.3 g/100 ml of potassium dichromate would have been obtained, just over the reported saturation point of 4.7 g/100 g at 0 C.

[Edited on 21-10-2012 by blogfast25]

KOH/air oxidation of chromite in aqueous suspension

tetrahedron - 24-10-2012 at 11:09

Green metallurgical processing of chromite

Quote:

A mechanical agitator was used at a stirring speed of
800 rpm to keep the slurry suspended during the leaching
experiment. A gas pipe with a distributor on one end
was fixed inside the bottom of the reactor and compressed
air was introduced through the pipe after prewetting.
Generally the gas flow rate was set at 0.4 m3/h.

Quote:

The desired amount of reagent KOH was first placed
in the reactor and mixed with a desired volume of distilled
water. The heating furnace was then switched on
and the system was heated. After most of the KOH was
dissolved, the agitator was turned on to improve the
mixing. When the temperature reached the preset value,
the chromite ore was fed into the aqueous solution and
the air compressor was started to provide the compressed
air.

with a ratio of KOH to chromite ore of 4:1 and after 2 hours of heating at 340°C, 80% conversion to chromate is reported.

edit. related article from the same authors: Oxidation decomposition of chromite ore in molten potassium hydroxide

[Edited on 24-10-2012 by tetrahedron]

blogfast25 - 25-10-2012 at 05:55

Very interesting find, tetra!

elementcollector1 - 25-10-2012 at 10:28

Chromite is blue-green? I think I might have a flask of that, my Cr(III) turned from emerald green to that color when baking soda in hot water was added.
More base turns it right back to chromate, though.

tetrahedron - 25-10-2012 at 11:36

Chromite is blue-green? I think I might have a flask of that, my Cr(III) turned from emerald green to that color when baking soda in hot water was added.
More base turns it right back to chromate, though.

i don't understand, by chromite the oxyanion of Cr(III) is meant. do you have other cations in there besides chromium and sodium?

Very interesting find, tetra!

glad you find it interesting =)

Plante:

Molten NaOH is dangerous overkill, as far as I'm concerned.

your method is surely good but i dont want to play with molten naoh at the moment.

i did it this way and i have to say i never felt any kind of danger. the lye easily melted in the crucible (a small stainless steel saucepan) on a propane stove, without splashing or other violent behavior

Don't forget that if you want dichromates salts you need to pass CO2 gas in the chromate solution and then add a satured solution of potassium chloride to your sodium chromate solution. Cool it down and filter precipitated potassium dichromate.

2Na2CrO4 + 2CO2 + H2O -) Na2Cr2O7 + 2NaHCO3

Sulphuric acid do the job too, but CO2 is cheaper and safer.

how would you generate the CO₂ without an acid? maybe coal combustion, but the CO coproduced might cause some reduction of the valuable Cr(VI).

Salmo - 25-10-2012 at 12:21

Hey everybody I just want to let you know that I'm proceeding with this thing even if university steal me almost all my time, tomorrow I will post you some new photos, I precipitated everything with NaOH and I washed the sludge with a lot of water to get ride of the sodium sulfate impurity, tomorrow I will add the NaOH to the sludge, I will heat everything first and than I will cool and add H2O2, I will then filter to obtain impure chromate that I will acidify and reduce with hydrogen peroxide to obtain chromium chloride that I will precipitate with sodium hydroxide.
I will than wash really well the precipitate, I will dry and weigh it and I will add the stoichiometric sodium hydroxide amount (heating and stirring well) plus hydrogen peroxide to obtain what I think will be almost pure sodium chromate.

Salmo - 26-10-2012 at 13:38

So long process.. anyway today I washed well the precipitate and heating the mass I added 20 grams (10grams excess) of sodium hydroxide (conc. solution) to the sludge, than I added 110ml of hydrogen peroxide ( I think it's enough, I had not time to do the calculations..) and I heated everything, tomorrow I will filter and I will go on with the procedure.

blogfast25 - 3-11-2012 at 07:56

Below is a photo of my potassium dichromate/KCl solution as obtained in the posts above, now with a small amount of K2Cr2O7 crystals:

Crystallisation (at the bottom of the 250 ml beaker) started a few days ago because it is now very cold in my lab. I helped it along by putting it in my lab fridge a couple more days.

Although the yield is clearly much lower than expected, it also shows I made enough dichromate for the solution to be saturated at about 5 C.

Purification of contaminated chromate

weiming1998 - 4-11-2012 at 03:01

A lot of the time, crystallisation of the crude chromate/dichromate will yield large amounts of salt, among other substances, that makes the crystals look almost white. I have found a method to purify such contaminated chromate.

1, Add excess CuSO4 to your crude chromate solution (please note that dichromate cannot be purified by this method). A very fine, lime-green precipitate will form. The CuSO4 doesn't have to be pure, it will work as long as it isn't adulterated by reducing agents or insolubles.
2, Heat the solution to make the precipitate clump together into less fine particles that can be filtered/decanted more easily.
3, Cool down solution, then filter/decant.
4, If you want really pure chromate, filter, dry and weigh the precipitate, then add a stoichiometric amount of NaOH to it. Add water and heat. If you don't care that much about the exact purity, then add water to the decanted/filtered precipitate, then while the solution is heated, slowly add NaOH solution to the mixture until the suspension of lime-green, insoluble particles turns black.
5, Whatever method you used, the solution will now be full of copper (II) oxide. Filter that out. Now boil down/evaporate the solution. This time, the resulting crystals will be bright yellow. The copper (II) oxide can be disposed, kept or recycled into CuSO4.

blogfast25 - 4-11-2012 at 06:56

This solubility products table:

http://www.csudh.edu/oliver/chemdata/data-ksp.htm

… has Ks = 3.6 x 10<sup>-6</sup> and 2.2 x 10<sup>-20</sup> for the respective solubility products of CuCrO4 and Cu(OH)2. That would fit what you claim quite nicely.

But it’s a fairly expensive way of separating, since as you obtain somewhat useless CuO as a by-product.

It could also be used to purify dichromates, by alkalising them slightly: even at neutrality the equilibrium dichromate < === > chromate will have been pushed sufficiently to the right for CuCrO4 to form.

[Edited on 4-11-2012 by blogfast25]

blogfast25 - 4-11-2012 at 07:19

Molten NaOH is dangerous overkill, as far as I'm concerned.[/rquote]

your method is surely good but i dont want to play with molten naoh at the moment.

Tetra, could you describe what you did more accurately?

For CO2 you'll need a generator (limestone + HCl e.g.) or bottled CO2 or soda water + acid. Something like that...

[Edited on 4-11-2012 by blogfast25]

tetrahedron - 4-11-2012 at 10:07

Tetra, could you describe what you did more accurately?

nothing special, i was referring to my attempt at the method suggested by plante:

i used stoichiometric amounts according to the reaction

Cr₂O₃ + 3KNO₃ + 4KOH ---> 2K₂CrO₄ + 3KNO₂ + 2H₂O

i don't have the numbers with me but i used about 10% excess pottery grade Cr₂O₃ and technical or better KNO₃ powder and 85% KOH pellets. the mix was stirred on the flame (low setting) until molten, then stirred occasionally while left mostly unattended for a few hours. better stirring would have probably promoted air oxidation and improved the yield. when i decided to stop the heating, it didn't look like much had happened. i let it cool, then leached the solid with distilled water and kept both phases with the prospect of quantitating the chromate formed as well as recycling the unreacted Cr₂O₃. hope this answers your questions, otherwise just ask specifically =)

i find the electrolytic methods more interesting and original so i'm concentrating my efforts on those. i tried the molten route for comparison.

For CO2 you'll need a generator (limestone + HCl e.g.) or bottled CO2 or soda water + acid. Something like that...

agreed..in that case why not use the acid to convert the chromate to dichromate directly? as to sodastream bottles, they are stoichiometrically in the same price range as hydrochloric or sulfuric acid. that's why i found plante's suggestion to 'just use CO₂' excessively nonchalant.

elementcollector1 - 4-11-2012 at 10:59

Quote: Originally posted by weiming1998

A lot of the time, crystallisation of the crude chromate/dichromate will yield large amounts of salt, among other substances, that makes the crystals look almost white. I have found a method to purify such contaminated chromate.

1, Add excess CuSO4 to your crude chromate solution (please note that dichromate cannot be purified by this method). A very fine, lime-green precipitate will form. The CuSO4 doesn't have to be pure, it will work as long as it isn't adulterated by reducing agents or insolubles.
2, Heat the solution to make the precipitate clump together into less fine particles that can be filtered/decanted more easily.
3, Cool down solution, then filter/decant.
4, If you want really pure chromate, filter, dry and weigh the precipitate, then add a stoichiometric amount of NaOH to it. Add water and heat. If you don't care that much about the exact purity, then add water to the decanted/filtered precipitate, then while the solution is heated, slowly add NaOH solution to the mixture until the suspension of lime-green, insoluble particles turns black.
5, Whatever method you used, the solution will now be full of copper (II) oxide. Filter that out. Now boil down/evaporate the solution. This time, the resulting crystals will be bright yellow. The copper (II) oxide can be disposed, kept or recycled into CuSO4.

Well, this gives me an excuse to make CuSO4. Are you sure this gets rid of all the other stuff?

(Reducing agents meaning acids?)

Is the CuSO4 in solution or on its own?

tetrahedron - 4-11-2012 at 11:44

this is a metathesis (precipitation) reaction based on the low solubility of copper chromate. any soluble non-reducing Cu(II) salt will do, as long as the anion won't precipitate any other cations in the crude chromate solution. e.g. CuCl₂ should work just as well. the Cu salt should be dissolved first. note that alkali sulfates are not as soluble as the corresponding chlorides (risk of co-precipitation).

Salmo - 4-11-2012 at 11:48

weiming this is what is was looking for!! CuSO4 is cheap and OTC, i will def try your method!

elementcollector1 - 4-11-2012 at 13:39

Made my own CuSO4 from sulfuric acid and copper electrochemically. It was still acidic, so I neutralized most of it with NaOH (never saw any Cu(OH)2, unfortunately) and added it to my chromate (heavily contaminated with NaCl, but still yellow). So, yellow solution + blue solution = green solution. No precipitate here, either. Have it cooling in the fridge.

blogfast25 - 4-11-2012 at 14:22

@tetrahedron:

Thanks.

Yes, I used direct acidification with HCl (KCl is more soluble than K2SO4, as Tim pointed out).

elementcollector1 - 4-11-2012 at 21:29

I don't think that copper precipitation method works for some reason. Unless my chromate or sulfate wasn't concentrated enough, I observed no precipitation at all. Someone want to try with concentrated, pure reagents and see if it works for them?

blogfast25 - 5-11-2012 at 14:16

Someone want to try with concentrated, pure reagents and see if it works for them?

That's altready been done by Weiming (see up). There's no reason at all to doubt his results, especially since theory predicts that such a method will work.

Your chromate solution is likely to be very dilute: after all you claim to make chromate without using an oxidiser (something which theory predicts to be impossible).

Be careful with making unsubstantiated claims...

[Edited on 5-11-2012 by blogfast25]

elementcollector1 - 5-11-2012 at 15:50

I made chromate with bleach and electrolysis of stainless steel, which I'm sure has been well documented thus far.

The dilution problem is likely to be correct, though.

tetrahedron - 5-11-2012 at 17:19

thank you all for the updates. i made some quick qualitative tests based on weiming's method of my three attempts at synthesizing chromate.

the CuSO₄ solution employed in all three tests:

1.
the first test appears to confirm the electrochemical oxidation of Cr(III) to Cr(VI) in NaOH/NaCl(aq) with inert anode as i described on page 3 of this thread:

yesterday i took some of the sludge that i got when "electrodissolving" stainless steel in a Cl^- containing electrolyte and poured it into a fresh solution of NaOH and NaCl (about 1 spoonful each in 400ml tap water), which i then electrolyzed for a couple hours with a 5mm diameter graphite "gouging rod" anode, initially @ ~7V, soon reduced to 4.4V & 3A in order to reduce the loss of chlorine.

each set of pics shows the sample of filtered product before and after the addition of the CuSO₄. note that in none of the tests did i neutralize the electrolyte beforehand. immediately a lime green precipitate formed:

2.
the second set refers to plante's method of electrolyzing Na₂CO₃(aq) with a stainless steel anode:

Only electrolysing S.S at 20V in Na2CO3 electrolyte. Sodium chromate go in to solution but iron and nickel don't. It take a large amount of time to get a product but it work really well.

clearly the basic Cu²⁺ precipitate completely overshadows any lime green (if present at all):

3.
the third and last set refers to the tried and true molten salt method: