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Author: Subject: Chromates from stainless steel! How does it work?
bluamine
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[*] posted on 2-5-2017 at 02:34
Chromates from stainless steel! How does it work?


Hi everyone!
I have watched the videos on the links below, which are about production of Na2CrO4 from stainless steel. I didn't realize the final step (oxidation with hypochlorite)
In other words, I would like to understand what is the equation of reaction between CrCO3 & NaOCl
https://youtu.be/F_W-IyUTM5M
https://youtu.be/3uzTNjuUyyk
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PirateDocBrown
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[*] posted on 2-5-2017 at 03:03


I would use Ca(OCl)2 if I were you. The calcium could then be precipitated by sulfate later.
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JJay
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[*] posted on 2-5-2017 at 03:23


The chromium carbonate first decomposes to chromium (iii) hydroxide, releasing carbon dioxide.

After that, the bleach oxidizes the chromium (iii) hydroxide to chromic acid, which reacts with sodium hydroxide in the bleach to form sodium chromate:

2 Cr(OH)3 + 3 NaOCl + 4 NaOH → 2 Na2CrO4 + 3 NaCl + 5 H2O




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bluamine
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[*] posted on 2-5-2017 at 03:31


Quote: Originally posted by JJay  
The chromium carbonate first decomposes to chromium (iii) hydroxide, releasing carbon dioxide.

After that, the bleach oxidizes the chromium (iii) hydroxide to chromic acid, which reacts with sodium hydroxide in the bleach to form sodium chromate:

2 Cr(OH)3 + 3 NaOCl + 4 NaOH → 2 Na2CrO4 + 3 NaCl + 5 H2O

Honestly I don't think household bleach contains that much of NaOH! Despite it's made by reacting its aqueous solution with chlorine gas
I think we should add some NaOH then

[Edited on 2-5-2017 by bluamine]

[Edited on 2-5-2017 by bluamine]
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bluamine
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[*] posted on 2-5-2017 at 03:35


Another thing to say is that the sodium chloride must be separated if we are planning to transform chromate to dichromate since chromate would react with sulfuric acid to form chromyl chloride (in the presence of NaCl)
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JJay
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[*] posted on 2-5-2017 at 03:52


The sodium chloride is actually pretty easy to get rid of.

I was actually thinking about attempting this on a 2M scale with calcium hypochlorite in the next few days. I think I might try adding some lime.

With all those metal hydroxides in the mix, you end up getting some basic chlorides and the sodium hydroxide regenerates, but last time I tried it, I did generate enough chlorine gas (or perhaps chlorine oxides) to frighten me into putting the flask in a bucket and taking it outside.

Wikipedia currently says that the sodium hydroxide concentrations of household bleach are extremely low (0.01–0.05%), but I think that is an error.

Edit: Chlorox doesn't mention the sodium hydroxide concentration in their bleach, but this brand says 0.3-5%: http://www.sunbeltchemicals.com/sites/default/files/product/...

I'm guessing that is probably typical of household bleach.

[Edited on 2-5-2017 by JJay]




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bluamine
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[*] posted on 2-5-2017 at 04:04


Quote: Originally posted by JJay  
The sodium chloride is actually pretty easy to get rid of.

I was actually thinking about attempting this on a 2M scale with calcium hypochlorite in the next few days. I think I might try adding some lime
.

Quote: Originally posted by JJay  
The sodium chloride is actually pretty easy to get rid of.

I don't think it's that easy, since chromate may oxidize any organic solvents if added to the mixture in order to precipitate NaCl out of solution!

Quote: Originally posted by JJay  
With all those metal hydroxides in the mix, you end up getting some basic chlorides and the sodium hydroxide regenerates, but last time I tried it, I did generate enough chlorine gas (or perhaps chlorine oxides) to frighten me into putting the flask in a bucket and taking it outside.

I would try to precipitate other metal chlorides first! Idk what are the best solvents which can be used for this purpose though

Quote: Originally posted by JJay  

Wikipedia currently says that the sodium hydroxide concentrations of household bleach are extremely low (0.01–0.05%), but I think that is an error.

I don't think it's very wrong!
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[*] posted on 2-5-2017 at 04:19


The metal chlorides are basic chlorides and are not very soluble, so no solvent is necessary.

If you boil the solution down, the sodium chloride crystallizes out cleanly. I've done it, and I just ran across this yesterday, which describes doing exactly that before adding a potassium salt: https://erowid.org/archive/rhodium/chemistry/potassium.dichr... It wouldn't hurt to check the product for chlorides, of course.

Sodium carbonate does not crystallize out nearly as cleanly from chromate solutions; it instead forms flakes that pull chromates out of solution, making a big mess.





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[*] posted on 2-5-2017 at 06:52


There's actually no such thing as chromium carbonate. Chromium(III) solutions precipitate as either the hydroxide or a mix of basic salts when treated with sodium (bi)carbonate solutions.

I truly don't know what the exact reaction here is, but pretending everything were nice neat hydroxides, My guess would be this:

2 Cr(OH)3 + 5 NaOCl = 2 Na2CrO4 + NaCl + 2 Cl2 + 3 H2O

Chlorine. Lots of chlorine seems to be generated every time I carry this out, despite completely neutralizing any acids leftover from the SS dissolution process. Industrial processes take advantage of chromium(III) oxide's amphoteric nature by leaching it out of ores as sodium chromate, using a base. So it might not be that outlandish to suggest that chromium(III) hydroxide behaves as an acid towards the bleach and liberates hypochlorous acid. Any way I try to look at writing out a reaction for this seems to suggest excess base added to household bleach would be a good idea.

Theory aside, my best results with this procedure have come from converting the chromate to dichromate with careful addition of acid (using too much hydrochloric is bad news, as it is reducing and can destroy the dichromate species; see Chem Player's video on potassium chlorochromate).
You can then add potassium chloride and evaporate or chill to crystallize out potassium dichromate.

[Edited on 5-2-2017 by Amos]
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[*] posted on 2-5-2017 at 09:43


It's tricky to know exactly how much bleaching powder to use to get a good yield because bleaching powder is almost always of questionable purity, and there are side reactions with the chromium compounds and the other metal hydroxides.



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[*] posted on 4-5-2017 at 23:16


I had planned on doing this yesterday but got delayed. Right now I'm going to attempt to oxidize about 367 grams of dissolved silverware. Not all of it was 18/8... there were a couple of items in the mix, including a mixer paddle and a spoon scoop that appear to be chrome plated. So the silverware that dissolved might be slightly richer in chromium than would be expected.

Assuming 18/8 proportions, there should be 66.11 grams of dissolved chromium (1.271 mol), 29.38 grams dissolved nickel (0.5006 mol), and 271.79 (4.867 mol) grams of iron.

Assuming that the chromium oxidizes to CrO3, the nickel oxidizes to Ni2O3, and the iron oxidizes to Fe2O3, and that each is in its hydroxide form, it looks like 912.9 grams of 99% Ca(OCl)2 are required. When I did this previously, I figured that chromium would tend to oxidize more easily than the other compounds and used about 2/3 the stoichiometric amount, but the yield was only about 50%. Some chlorine gas was produced, but in retrospect, there wasn't really that much of it. If I use the full amount of bleaching powder, I expect a lot more chlorine... I think chlorine production will be greatly reduced and perhaps the yield increased if I first mix in some lime... and it looks like 94.17 grams (1.271 mol) are required.

---

After adding about 100 grams food grade calcium hydroxide and about 100 grams of "99%" bleaching powder ("68% chlorine"), I can see that a reaction is taking place as the bleaching powder dissolves, turning the green sludge orange, and it is fizzing and giving off a gas, but that gas doesn't seem to be anything noxious. I guess it might be carbon dioxide from iron carbonate oxidizing. I can't smell any chlorine in the reaction mixture at all. I'm tempted to just dump the entire two pounds into the bucket, but I should probably at least take it outside first if I am going to do that.

---

After adding two entire pounds of bleaching powder (outside), I can barely smell any chlorine coming from the bucket. I am going to let it react for a couple hours or so, and then, if it doesn't smell like the business end of a water treatment plant, I will bring the bucket inside. Then I'll let it sit around for a few days to settle while I try to figure out how to filter the reaction mixture....

---

I can definitely smell some chlorine in it but not that much, and it's still reacting, but I brought it inside and it's sitting in a room with a window fan.

---

After about 20 minutes, I decided that there was too much chlorine coming off of it for comfort and put it back outside to finish reacting.

[Edited on 5-5-2017 by JJay]




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[*] posted on 5-5-2017 at 15:08


I let it continue reacting for several hours and observed that it had only a little bit of chlorine smell and brought it inside. I also observed that stirring caused it to emit a sharp and distinctive halogen smell (all halogens smell pretty much the same - if you know the smell of iodine, you know the smell of chlorine and bromine). This was actually a bit different than the "chlorine" smell it was giving off previously, which smelled basically like bleaching powder. I had thought a couple of weeks ago that this smell was probably chlorine dioxide, and that seems to make sense in light of the sharp contrast with the smell that is coming off of it now (which actually smells like chlorine).

I have very strong ventilation running, but I might have to take the bucket back outside.

---

I stirred it a few times with a paddle and now I have it sitting on a magnetic stirrer with a large stir rod. It's kicking up an impressive vortex considering that the mixture is around 10L. There is a slight chlorine scent being given off. The mixture seems to consist of a yellow solution containing a lot of extremely fine red powder.

[Edited on 6-5-2017 by JJay]




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