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Author: Subject: Reduction of Hexavalent Chromium
elementcollector1
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[*] posted on 22-8-2012 at 20:02
Reduction of Hexavalent Chromium


Before I re-attempt to isolate chromium from an ungodly amount of scrap stainless steel that I need to get rid of, how would I reduce hex-chrome (let the gasps and danger warnings begin!) to good ol' tri-chrome? For safety reasons, you know.



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[*] posted on 22-8-2012 at 22:31


Hexavalent chromium is easily reduced by many reductors. First, assure that the chromium comes in aqueous solution in some form. Next, acidify this solution and then add a suitable reductor. Sulfites or (meta)bisulfites are really good. They instantaneously reduce hexavalent chromium to trivalent chromium. Ethanol also does the job, but with this, the reduction takes some time and it is best to heat the liquid somewhat to speed up things.



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[*] posted on 24-8-2012 at 11:59


And would the addition of more ethanol precipitate out the Cr(III) compound? One organochem professor mentioned to me that the addition of alcohol to an aqueous solution messes with the solubility, precipitating out the compound.
Where do I get metabisulfites?
Does the ethanol need to be free of water for best effect?




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[*] posted on 24-8-2012 at 21:52


Quote: Originally posted by elementcollector1  
And would the addition of more ethanol precipitate out the Cr(III) compound? One organochem professor mentioned to me that the addition of alcohol to an aqueous solution messes with the solubility, precipitating out the compound.
Where do I get metabisulfites?
Does the ethanol need to be free of water for best effect?


There is no need for anhydrous ethanol/metabisulfite. If you want to convert Cr(VI) to Cr (III), then just do the following :

1, Dissolve your Cr(VI) source in water and add sulfuric acid to catalyse
2, Add excess ethanol/isopropanol, whatever is available (any organics that has a hydroxyl functional group (except for tertiary alcohols) will work for this reduction).
3, Boil down solution to get rid of volatile organics and water.
4, You have now got a Cr (III) salt.
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[*] posted on 25-8-2012 at 14:23


Thanks! Concentration of the sulfuric acid? (I have this terrible pink 10% stuff...)



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[*] posted on 25-8-2012 at 16:00


Quote: Originally posted by elementcollector1  
Thanks! Concentration of the sulfuric acid? (I have this terrible pink 10% stuff...)


Any will do. Sulfuric acid is added just so there are more H3O+ in the water. Sodium bisulfate solution from pool stores will even work.
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[*] posted on 26-8-2012 at 01:38


With a bit of luck the Cr(VII) will destroy whatever the pink colour is too.
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[*] posted on 26-8-2012 at 08:49


The electrolysis is going inordinately slow, which is odd, because it's a one-cell, NaCl-saturated solution with SS as both electrodes. What do you think is happening?
(Also, as of the first hour, there is a slight yellow tinge to the cell, could be chlorine or chromate at this point.)

EDIT: Hour 2: Greenish-grayish dark precipitate formed. Chromium hydroxide or Ferrous hydroxide?

[Edited on 26-8-2012 by elementcollector1]




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[*] posted on 26-8-2012 at 11:51


Quote: Originally posted by elementcollector1  
The electrolysis is going inordinately slow, which is odd, because it's a one-cell, NaCl-saturated solution with SS as both electrodes. What do you think is happening?
(Also, as of the first hour, there is a slight yellow tinge to the cell, could be chlorine or chromate at this point.)

EDIT: Hour 2: Greenish-grayish dark precipitate formed. Chromium hydroxide or Ferrous hydroxide?

[Edited on 26-8-2012 by elementcollector1]


What electrolysis are you referring to?

Another great reducing agent of Cr (VI) (as dichromate Cr2O7 (2-)), often overlooked, is actually H2O2:

Cr2O7(2-) + +14 H+ + 6 e → 2 Cr3+ + 7 H2O
3 x [H2O2 → O2 + 2 H+ + 2 e]

Peroxide oxidises Cr3+ to chromate in alkaline conditions but dichromate in acid conditions oxidises peroxide…




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[*] posted on 26-8-2012 at 12:19


One-cell, NaCl-saturated electrolysis of 2 stainless steel electrodes. Makes chromate! (well, you need to add a bit of HCl to get the bubbles going, but still.)
Wasn't there an unstable Cr (V) peroxocomplex or some such?




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[*] posted on 26-8-2012 at 22:21


Yes, there are many different peroxo complexes of chromium(VI) and of chromium(V). The most stable is the dark brown salt K3Cr(O2)4, potassium tetraperoxochromate(V). There also are blue complexes, the unstable CrO(O2)2, which is a chromium(VI) complex and the red/brown chromium(IV) complex Cr(NH3)2(O2)2. There are many more complexes, but these three are the most common ones. I prepared the K3CrO8 complex and still have that around, it is stable, but explodes when heated above a flame. I also made Cr(NH3)2(O2)2, but this complex deteriorates on storage. One month after initial preparation it already has mostly decomposed.



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[*] posted on 27-8-2012 at 04:27


To reduce Cr(VI) to Cr(III) with H2O2 these peroxo complexes are easily avoided.



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[*] posted on 27-8-2012 at 04:30


Quote: Originally posted by elementcollector1  
One-cell, NaCl-saturated electrolysis of 2 stainless steel electrodes. Makes chromate! (well, you need to add a bit of HCl to get the bubbles going, but still.)
Wasn't there an unstable Cr (V) peroxocomplex or some such?


And this actually works? You have evidence/proof that chromate has been formed in significant quantities?




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[*] posted on 27-8-2012 at 05:23


Quote: Originally posted by blogfast25  
Quote: Originally posted by elementcollector1  
One-cell, NaCl-saturated electrolysis of 2 stainless steel electrodes. Makes chromate! (well, you need to add a bit of HCl to get the bubbles going, but still.)
Wasn't there an unstable Cr (V) peroxocomplex or some such?


And this actually works? You have evidence/proof that chromate has been formed in significant quantities?


This would indeed work in theory. The anode of the stainless steel electrode will slowly be worn away by the chlorine produced to form FeCl3, along with small amounts of CrCl3, which reacts with the excess NaOH to form both Fe2O3.H2O and Cr(OH)3, which is the method of producing iron oxide by electrolysis (with iron anodes). The chromium hydroxide/oxide will be oxidized by the hypochlorite produced on site to form chromium (VI) compounds, primarily sodium chromate. Here's a reference to that http://www.sciencedirect.com/science/article/pii/S0020169306...

But the question is, why not dissolve the stainless steel in an acid first, then oxidize by hypochlorite to Cr(VI) compounds? The formation of ferrates can be prevented by dropping the pH after the reaction, so that it decomposes. Better yet, take the boiled down solids produced by dissolving stainless steel in acids and oxidize in air by molten NaOH.
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[*] posted on 27-8-2012 at 06:38


Are you using nickel stainless steel (18/10) or nickel-free stainless steel (18/0)? I used the latter and it dissolves very easily and quickly in a one-cell NaCl electrolysis apparatus. (If the stainless steel is bluish-tinted and ferromagnetic, it is 18/0. If it is golden tinted and not magnetic, it is 18/10.)



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[*] posted on 27-8-2012 at 10:53


Weiming:

Yes, if obtaining chromium compounds from FeCr alloys is the goal here, then dissolving the lot in HCl, precipitating all as hydroxides, then alkaline oxidising the chromium (III) to soluble chromate (VI), is far easier than electrolysis, IMHO.




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[*] posted on 27-8-2012 at 12:45


I don't know about that, it seems to be going well on my end. (although the chemical method was faster, I don't have any NaOH on hand. Would bleach work?)



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[*] posted on 27-8-2012 at 12:55


Probably but I wouldn't really recommend it.



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[*] posted on 27-8-2012 at 14:34


Why, chlorine?



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[*] posted on 28-8-2012 at 01:25


Quote: Originally posted by elementcollector1  
Why, chlorine?


An aqueous synthesis of chromates with sodium hypochlorite will probably work, as I said before, but the amount of impurities produced are high, the overall process isn't very efficient, and lots of bleach are needed for a small amount of chromate.

Why not make some NaOH? Get a bag of slaked lime from where you get fertilisers, and react that with boiling sodium bicarbonate solution, then filter.
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[*] posted on 28-8-2012 at 04:54


Quote: Originally posted by weiming1998  
Why not make some NaOH?


Why not buy some? Cheap as chips. Or KOH from the Biodiesel People...

[Edited on 28-8-2012 by blogfast25]




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[*] posted on 28-8-2012 at 09:36


Thats totally not the way they recycle SS, at best you could make lesser ferrochrome alloys by adding melts of ss into a cast of molten iron.
Why not just buy the ore? You gonna spend much more on this multistep load of waste of chemicals first redissolving the already neutral atom of chromium to reduce it again afterwards. Check on the prices for ore exports from pakistan or something its ridiculous cheap, also, you can easily separate iron from Cr in oxide form.
A whole ton with around 40% chromite goes for 200 bucks if I'm correct LOL


Also the procedure in aqueous solution used to reduce Cr(VI) to Cr(III) involves a lot of messy ions in solution, like sodium, which are problematic to separate afterwards. I would suggest making an ammonium based salt of dichromate prior to putting it to react.

Na2Cr2O7 + H2O + 2Ca(OH)2 --> CaCrO4 + 2NaOH +2H(+) (something went up wrong here)
Boil
CaCrO4 will ppt, filter
Cool, dilute it again
Add (NH4)2SO4, wait until dissolves.
Collect CaSO4 ppt, you now have (NH4)Cr2O7
Put this to oxydise ethyl alcohol by the method described by weiming 1998, use nitric acid as the catalyst.
Now you have a mess of salts, but all of them can be decomposed by heat leaving pure chromium compounds behind.
I will just do this now see what happens.


[Edited on 8-28-2012 by Poppy]

[Edited on 8-28-2012 by Poppy]




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[*] posted on 28-8-2012 at 10:59


The ammonium dichromate decomposition method was discussed in several other threads, and is NOT a good method of getting pure chromium (III) oxide.
Besides, if I reduced the Cr(VI) to Cr(III), and then added a bit of ammonia, I get instant chromium hydroxide. Eventually, I'll get a large amount of this, and then heat it to decompose to pure Cr2O3 and water.




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[*] posted on 28-8-2012 at 11:46


Rather, tell your mate the aluminium foil has been dismissed. Waste all the aluminium in the following reaction:

Cr2O7(2-) + 14H+ + 2Al --> 2 Cr(3+) + 7H2O + 2Al(3+)




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[*] posted on 28-8-2012 at 15:16


Wouldn't ethanol be better as a reducing agent? It'll get destroyed when I boil the resulting Cr (III) solution down...



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