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Author: Subject: Reduction of Hexavalent Chromium
blogfast25
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[*] posted on 10-9-2012 at 09:17


Quote: Originally posted by elementcollector1  
Excellent! That second thread was the one I was talking about.
Anyway, some of that strange metallic float disappeared, but it still remains, oddly enough, in the already filtered flask.


The 'metallic' float you're referring to is something that can be seen quite often on the solutions of some metals. It's likely to be caused by thin layer diffraction of a thin layer of insoluble metal oxides of hydroxide. I've seen it form on solutions of Fe2+ exposed to air, for instance, and it does have a bit of a metallic sheen. But what you're seeing isn't metal. And when this film acquires thickness it will eventually sink.

[Edited on 10-9-2012 by blogfast25]




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[*] posted on 11-9-2012 at 07:19


Aw. I was hoping that chromium had somehow formed and that I could just scoop it up.
I really couldn't tell it wasn't oxides, it looked as shiny as...well... chrome. :P




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[*] posted on 11-9-2012 at 09:22


Quote: Originally posted by elementcollector1  
Aw. I was hoping that chromium had somehow formed and that I could just scoop it up.
I really couldn't tell it wasn't oxides, it looked as shiny as...well... chrome. :P


'Spontaneous reduction' of chrome, eh? That would have been convenient, just not very likely... ;)




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[*] posted on 11-9-2012 at 11:03


Well, at least my solution is a golden yellow. Not quite Gatorade yellow the way I've heard it's supposed to be, but we'll see.
By the way, what should I do to get the solid chromate? I'm afraid to boil it because of the hex-chrome being released as fumes, evaporating takes forever, but I could dessicate it. What do you think?
(I have roughly 1.5-2 liters of the raw liquid, by the way.)




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[*] posted on 12-9-2012 at 06:53


Quote: Originally posted by elementcollector1  
Well, at least my solution is a golden yellow. Not quite Gatorade yellow the way I've heard it's supposed to be, but we'll see.
By the way, what should I do to get the solid chromate? I'm afraid to boil it because of the hex-chrome being released as fumes, evaporating takes forever, but I could dessicate it. What do you think?
(I have roughly 1.5-2 liters of the raw liquid, by the way.)


I thought your purpose was to obtain Cr2O3 to obtain chromium metal? In which case just reduce the lot to Cr3+, then precipitate with soda, as Cr(OH)3 hydrate. Filter, wash and semi-calcine.

Dichromate solutions can be safeky boiled in though, although 2 L is a lot of boiling! Just make sure the boiler/container is partly covered to avoid droplets of solution getting airborne.

Dessicating would take a ton of dessicant and an eternity of time.

[Edited on 12-9-2012 by blogfast25]




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[*] posted on 12-9-2012 at 09:27


Alright. I just want to make the solution as concentrated as possible before reducing it, so if I boil it to dryness, weigh it, and add enough water to dissolve most of it (supersaturated solution), that would make the most Cr(OH)3 in one go.
By soda, do you mean baking soda?




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[*] posted on 12-9-2012 at 10:19


It’d be quicker and easier to reduce it now. Then add soda (sodium carbonate, not baking soda), which has just the right alkalinity to precipitate Cr(OH)3 although you may want to neutralise much of the acidity with NaOH first (so as not to waste too much soda). Allow to stand overnight (collecting the hydroxide at the bottom), then filter and wash.

Or you could concentrate it by reducing the volume by half first.

But to extract significant amounts of Cr from stainless steel requires quite a bit of SS, at about 12 % Cr only…


[Edited on 12-9-2012 by blogfast25]




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[*] posted on 12-9-2012 at 16:41


Can I make the carbonate from the bicarbonate? I thought it might be possible by boiling a solution of bicarbonate to dryness, or some such.
Alternatively, I think I have some 'washing soda' somewhere...




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[*] posted on 13-9-2012 at 05:38


Quote: Originally posted by elementcollector1  
Can I make the carbonate from the bicarbonate? I thought it might be possible by boiling a solution of bicarbonate to dryness, or some such.
Alternatively, I think I have some 'washing soda' somewhere...


Bake the baking powder in the oven for about 2 h @ 200 C, it converts completely to sodium carbonate:

2 NaHCO3(s) == > Na2CO3(s) + H2O(g) + CO2(g) (that's how baking soda works!)

You may want to check the quality of the obtained soda by dissolving some in water: if turbid (anti-caking agents!), filter.

But soda is readily available from hardware stores as 'washing soda'. It too may require filtering prior to use as a chemical reagent.

[Edited on 13-9-2012 by blogfast25]




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[*] posted on 13-9-2012 at 06:30


In addition to hardware stores, washing soda is also available in some grocery stores and supermarkets that sell laundry supplies.
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[*] posted on 13-9-2012 at 15:19


Well, my 'chromate' evaporates to white crystals. I'm assuming the bleach is playing a part here. This shouldn't be a problem, as all that bleach should disappear upon addition of acid (and then ethanol, and then soda).



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[*] posted on 14-9-2012 at 05:41


Quote: Originally posted by elementcollector1  
Well, my 'chromate' evaporates to white crystals. I'm assuming the bleach is playing a part here. This shouldn't be a problem, as all that bleach should disappear upon addition of acid (and then ethanol, and then soda).


Your bleach is no longer bleach (sodium hypochlorite) but sodium chloride. If there's a lot of it, it will definitely completely obscure the dichromate.

Redissolve in a small amount of water, reduce with alcohol + acid, then add soda solution till fizzing starts and Cr(OH)3 has precipitated. Allow the hydroxide to settle, the supernatant liquor should then be clear and colourless. Filter, wash and dry.

[Edited on 14-9-2012 by blogfast25]




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[*] posted on 14-9-2012 at 15:36


Well, I boiled the liquid down to a more concentrated form (100mL to I'm guessing about 20-30 mL). The precipitate crystals are a pale yellow-orange, while the supernatant liquid is a deep yellow-orange. Unfortunately, I don't have any acid on me, and I'm restocking tomorrow (Hello McLendon's!) I'll try to get some pictures of the stuff through the process.



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[*] posted on 15-9-2012 at 05:56


A word about precipitating Cr(OH)3 with sodium carbonate.

Assuming you’ll carry out the reduction of the dichromate with alcohol in acid conditions, you’ll obtain the chromium as chromic cations (Cr(H2O)6(3+)). But in the trivalent state, Cr is amphoteric and also forms soluble chromite anions (blue/green) in strongly alkaline conditions:

Cr(H2O)6(3+)(aq) + 4 OH-(aq) → Cr(OH)4(-)(aq) + 6 H2O(l) (very simply put here)

That’s why sodium carbonate is a better alkali here than sodium hydroxide: it’s not as alkaline and the chance of Cr staying in solution as chromite is much smaller.

However, the acid reduction means there’s quite a bit of acid reserve in the reduced solution. To avoid wasting too much sodium carbonate, carefully use strong sodium hydroxide to neutralise most of the acid first, to a pH of about 4 – 5 – 6. Then complete the neutralisation/precipitation with sodium carbonate. This method avoids ‘overshooting’ neutrality too much and ending up with a chromite solution instead of a chromic hydroxide precipitate. Don’t use an excess sodium carbonate either: just use what’s needed to get to a pH of 8 – 9. That’ll happen very shortly after fizzing (CO2(g)) stops. Note that with all that neutralising your solution will heat up due to neutralisation enthalpy: always go slowly forward, especially the first time you do this…

After the precipitation/neutralisation the supernatant liquid should be clear and colourless: any green/blue hue would point to chromite in solution. In that case re-acidify and start again…




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[*] posted on 15-9-2012 at 12:52


I would not encourage pouring vessel contant as fast as possible involving concentrated conjugates at demonic speed like sulfuric acid and soda as to "make it work" without delay.
Instead, drip it very slowly as some clumps may form which surrounds its interior in a bubble like fashion thus enables for I would put 15% of the mess to be co-precipitated.
That means Na in your Cr which you can't filtrate. I.E: chormite forming locally.
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[*] posted on 15-9-2012 at 18:26


Wouldn't the Cr be the only thing precipitating?
I just got all my bases and acids restocked (all my base are belong to me), so I will give this a go tomorrow.
Still can't find washing soda for some reason... how long does baking soda take to decompose at 200 C? If I turn it up to 450 C, will it go faster? (In the realm of 30 minutes is good, I usually have to explain my way through putting random stuff in the oven).




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[*] posted on 16-9-2012 at 05:49


Re. Poppy’s remark, as with all precipitations it’s recommended to add the precipitating reagent quite slowly with constant stirring of the target solution, to avoid local high concentrations which can lead to co-precipitation of other substances (stuff that gets trapped in the precipitate's crystal lattice of water cloud), even local chromite ‘trapped’ in the Cr(OH)3 precipitate. It’s a general ‘rule’.

Temperature/reaction speed relation is well understood and follows from collision theory. It applies here just the same: at 450 C the decomposition of sodium bicarbonate to sodium carbonate should be over in about 30 min. You’ll probably have to grind down the obtained product because it will be a little more resistant to dissolution and filtering may be necessary to remove some bits that resist dissolving. I make anhydrous Na2CO3 for acid/base titrations regularly and have noticed that this ‘baked’ product can take some stirring and shaking to dissolve it completely. Of course you can also heat the solution to speed things up a bit.


[Edited on 16-9-2012 by blogfast25]




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[*] posted on 16-9-2012 at 11:34


It appears I added so much H2SO4 that the chromate was reduced anyway, so I'll just have to precipitate?
I have baking soda being nuked in the oven as we speak, but there are no physical signs of it decomposing earlier (other than minor bubbling and melting around the edges). Doesn't it release water and CO2 as gases? If so, shouldn't it be... doing something?
Well, anyway, this seems a viable route to pure chromium oxides. It was a bit longer than expected, but now that I have sodium hydroxide, that shouldn't be a problem.

EDIT: I had some leftover stained MnO2 on the baking dish that I couldn't get rid of, and either this appears to be spreading or my baking soda is decomposing into carbon. What is going on?

[Edited on 16-9-2012 by elementcollector1]




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[*] posted on 16-9-2012 at 13:41


You've been cleaning your oven lately? Could be fat and tar dripping from the oven walls and roof. lol
believe me, meat sauces have a well known tendency to explode and splash all around, no exception inside that burning cage called oven.
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[*] posted on 16-9-2012 at 18:18


Um, no? Whatever it was, it appeared to have ruined the entire batch. Got practically no Cr(OH)3 out of that 20mL of concentrate.
Perhaps I shall make a stop at the pottery department and get some Cr2O3 there. Any suggestions on purity?




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[*] posted on 17-9-2012 at 09:43


EC:

H2SO4 cannot reduce Cr(VI) all by itself. Perhaps you used peroxide for the oxidation? Left over peroxide will reduce Cr(VI) in a jiffy in acid conditions. Was your solution green/blue now?

When ‘baking’ something in an oven, ALWAYS, part cover it. The best way to verify if the bicarbonate has been converted to carbonate is to weigh before and after. The weight loss should correspond to the reaction described above. Also, sodium carbonate formed this way has a tendency to tick to glass, in my experience. I use silicone baking dishes for that kind of thing.

Pottery Cr2O3 is of unknown purity. Highly calcined it’s also difficult to dissolve in strong acids (and thus analyse). Pottery pigments often contain some silica but it’s really on a case-by-case basis.





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[*] posted on 17-9-2012 at 11:55


Darn! It was light green at first, then shifted to emerald green...
My baking dish was a ceramic plate that had previously been used for drying MnO2 'mud', and bore several stains pertaining to that.
So, for example, 50g of bicarbonate would decompose to 31g of carbonate? (Source: http://theodoregray.com/PeriodicTable/MSP/BalanceReactions)
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[*] posted on 21-9-2012 at 11:33


Quote: Originally posted by elementcollector1  
Darn! It was light green at first, then shifted to emerald green...
My baking dish was a ceramic plate that had previously been used for drying MnO2 'mud', and bore several stains pertaining to that.
So, for example, 50g of bicarbonate would decompose to 31g of carbonate? (Source: http://theodoregray.com/PeriodicTable/MSP/BalanceReactions)
^Most useful online tool, ever.


Yes, correct: 2 x 84 g gives 106 g, so 50 g gives 31.5 g.




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[*] posted on 16-10-2012 at 18:02


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?


that yellow brew that results from the electrolysis of stainless steel can be very tempting, although i haven't been able to detect chromate in it..more likely it's a mix of orange iron (hydr)oxides and green chromium (III) (hydr)oxides, net result = yellow.

i switched to a dilute NaCl electrolyte (something like 1 teaspoon/L; no HCl!) and the 18/10 anode (fork) dissolved within a couple hours leaving behind a fine dark brown sludge that i painstakingly washed by filtering through 2 coffee filters, then tried to dry over a gas burner (this resulted in a thick paste reminiscent of mortar; i gave up on trying to achieve complete dryness). the filtrate was very pale (i'll keep it for the next electrolysis).

Quote: Originally posted by elementcollector1  
Well, at least my solution is a golden yellow. Not quite Gatorade yellow the way I've heard it's supposed to be, but we'll see.


gatorade yellow is not a good description..look up chrome yellow, there's a difference

let us know how it goes =)
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[*] posted on 17-10-2012 at 15:46


Well, it reduced just fine, sulfuric acid and isopropyl really does the trick.
I'm still wondering about the sodium carbonate. I poured some boiling water over the bicarbonate, and it fizzled and bubbled as if decomposing (the gas was likely CO2). Is this pure sodium carbonate? What say you, blogfast25?
(I don't know, it still looks like Gatorade to me. Then again, there's always that 'named colors' dilemma with our eyes...)




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