Sciencemadness Discussion Board - Reduction of Hexavalent Chromium

Reduction of Hexavalent Chromium

elementcollector1 - 22-8-2012 at 20:02

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.

woelen - 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.

elementcollector1 - 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?

weiming1998 - 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.

elementcollector1 - 25-8-2012 at 14:23

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

weiming1998 - 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.

unionised - 26-8-2012 at 01:38

With a bit of luck the Cr(VII) will destroy whatever the pink colour is too.

elementcollector1 - 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]

blogfast25 - 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…

elementcollector1 - 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?

woelen - 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.

blogfast25 - 27-8-2012 at 04:27

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

blogfast25 - 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?

weiming1998 - 27-8-2012 at 05:23

Quote: Originally posted by blogfast25

Quote: Originally posted by elementcollector1

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.

LanthanumK - 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.)

blogfast25 - 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.

elementcollector1 - 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?)

blogfast25 - 27-8-2012 at 12:55

Probably but I wouldn't really recommend it.

elementcollector1 - 27-8-2012 at 14:34

Why, chlorine?

weiming1998 - 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.

blogfast25 - 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]

Poppy - 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]

elementcollector1 - 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.

Poppy - 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+)

elementcollector1 - 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...

Poppy - 28-8-2012 at 22:26

CrO3 when mixed with ethanol reacts vigorously to form Cr (II, III).
I got a small crucible I'll give it a try for sure making some of it.

Poppy - 30-8-2012 at 17:11

I tried producing the chromium trioxide by adding anhydrous sodium dichromate (dried in the oven: at first you may say "omg it caked", but as it slowly cools it easily releases from the glass) to 98% sulfuric acid. It dissolved, warming the vessel. I did not measure but I would put the temperature around 60°C, was just hot to touch, guess it could be heated even more au bain marie. Also realised the more excess sulfuric acid you add turns the mess into a better oxydising paste, no better than actual aqueous solution of both ingredients

. After cooling it becomes a dark red mud.
Analysing the reaction
H2SO4 + Na2Cr2O7 --> CrO3 + Na2SO4 + H2O
The problem obviously is separate the sulfuric acid and water from the rest of the contents. CrO3 decomposes at 251°C, while sulfuric acid boil at more than 300°C. So a vaccum pump mut be put to work as to easily get rid of the sulfuric acid.

Again I insist useing ammonium componds as the product of the dichromate because the procuct ammonium sulfate can evaporate too.

For the synsthesis of ammonium dichromate, do as follows, it has been succesfully tested:
(I) Na2Cr2O7 + 2 NaOH --> 2 Na2CrO4
(just add solid NaOH to conc. solution of dichromate)
(II) Na2CrO4 + CaCl2 --> CaCrO4 (ppt) + NaCl
Make a ten-fold excess solution of calcium chlorid so as to pour the conc. sodium chromate, this will render a lower degree of sodium contamination and also ensure reaction will dislocate as to precipitate CaCrO4 by means of common ion effeft. To this point CaCrO4 precipitates as a hydrate. Filter. Throughly dry in the oven, at 180°C. Now, the anhydrous salts may be washed with boiling water: very few of the now anhydrous CaCrO4 will dissolve, while all other salts will be taken away.
next
(III) 2CaCrO4 + 2(NH4)HSO4 --> (NH4)2Cr2O7 + 2CaSO4
Prepare a saturate solution of ammonium bisulfate at 80°C, pour the anhydrous CaCrO4 and stir to dissolve. This can be next diluted to facilitate filtration.
Now you have the precious ammonium dichromate!
:cool:

Attach: sulfuric acid vapour pressure graph

sulfuric acid vapour pressure.JPG - 30kB

[Edited on 8-31-2012 by Poppy]

elementcollector1 - 30-8-2012 at 17:41

Is there any OTC source for ammonium bisulfate?
Once I have the ammonium dichromate, what do I do with it?

EDIT: CrO3 -(197+ C)> Cr2O3 + O2
Why isn't this feasible?

[Edited on 31-8-2012 by elementcollector1]

[Edited on 31-8-2012 by elementcollector1]

Poppy - 30-8-2012 at 20:42

The ammonium dichromate carries the chromium content you dealing with all the time!!!

Its a bet you can achieve the high purity light green Cr2O3 and then experiment with CrO,
H3PO2 + 2 Cr2O3 → 4 CrO + H3PO4
Then check what wiki means about the decomposition of CrO at 300°c
Also, check the sticky on the preparation of elemental phosphorus!

Good luck!

elementcollector1 - 30-8-2012 at 21:12

I read somewhere on this forum (thread was complete with pictures) that the decomposition of ammonium dichromate does not produce pure Cr2O3, it instead produces a rather dark green powder ripe with impurities
(Ze search function fails me yet again, I can't find ze thread...)

Poppy - 30-8-2012 at 21:50

Yeah probably the thread is somewhere, couldn't find it either.
The decomposition method yields a low purity Cr2O3 powder for which there is no much description regarding its behaviour against acids.
I would dream of an equation where CrO desproportionates to Cr and some other chromium oxide!
like
3CrO --> Cr + Cr2O3 !!!!
This could then be purified useing CO to form carbonyl complexes out of the metallic chromium, leaving the oxides unreacted, while Cr(CO)6 might be solvated and extracted

[Edited on 8-31-2012 by Poppy]

elementcollector1 - 31-8-2012 at 13:40

Washing with distilled, boiling water would probably improve the quality of the chromium oxide, because that would remove all soluble compounds like unreacted dichromate, sodium/ammonium compounds, etc, leaving behind only a mix of chromium oxides (CrO, Cr2O3, and CrO2).
This would probably still work for a thermite.

I've seen several videos of the ammonium dichromate decomposition, and while it is insanely cool, I would probably prefer to just decompose the CrO3 (unless it's an explosion or something).
Is there a reaction for the decomposition of sodium dichromate?

Poppy - 31-8-2012 at 14:16

No no, sodium holds everything together for longer. It turns into a melt, and stays like that, specially the dihydrate.

elementcollector1 - 31-8-2012 at 22:04

Hmm. Well, how do I turn sodium chromate to pure sodium dichromate? Then I can start the conversion to CrO3, and probably decompose from there.

blogfast25 - 1-9-2012 at 05:22

Forget about decomposing ammonium dichomate thermally, it's not great for pure Cr2O3.

Instead dissolve it in water, add some acid and add (slowly!) methylated spirits, then gently heat if necessary. The alscohol reduces the Cr (VI) to Cr (III) (green/blue). Now alkalise the solution with ammonia and Cr(OH)3.nH2O precipitates. Filter this off, wash with copious amounts of water and dry/calcine to Cr2O3.

Instead of alcohol, hydrogen peroxide can also be used.

[Edited on 1-9-2012 by blogfast25]

elementcollector1 - 1-9-2012 at 21:54

Er, has anyone heard what I've said? About CrO3 decomposing to Cr2O3? I've gotten no feedback on that so far, and it would be great if I could just skip most of the steps to that part.

Blogfast, what if I used sodium dichromate instead?

blogfast25 - 2-9-2012 at 05:16

Quote: Originally posted by elementcollector1

Er, has anyone heard what I've said? About CrO3 decomposing to Cr2O3? I've gotten no feedback on that so far, and it would be great if I could just skip most of the steps to that part.

Blogfast, what if I used sodium dichromate instead?

CrO3? Basically I'd forget about that if I were you.

All dichromates reduce alcohols and H2O2: it's a function of the anion, not the accompanying cation (K, NH4, Na,...)

elementcollector1 - 2-9-2012 at 19:42

Aww. I thought that might work, given how there would be no extra stuff in CrO3, just more oxygen.
Well, I switched to the chemical method a while back,and have a mix of my stainless steel chlorides and bleach reacting and filtering away. Last I saw, it was just beginning to turn yellow.

Poppy - 2-9-2012 at 21:19

Could you describe what's that you performing with the ss you have, becaue a bunch of SS is sometimes difficult to dissolve

elementcollector1 - 2-9-2012 at 21:22

1. Dissolved the stainless in a bunch of concentrated HCl. After a while, the solution darkened from clear to emerald green to opaque green. No reflux was needed, just some time.
2. Added bleach. Ferric hydroxide precipitated out in large quantities, this is busy being filtered out while the bleach additionally reacts with the chromium (III) hydroxide precipitated to form sodium chromate.

keep the acid

Poppy - 2-9-2012 at 21:27

If your HCl goes weak you should attach a HCl generator lined into the batch vessel, the use of a non expensive silicon rubber tube should suit for a couple runs

elementcollector1 - 6-9-2012 at 19:37

Well, I got what I thought was a large quantity of sodium chromate, but upon adding acid it did not change color to the orange dichromate, so... Any thoughts?

Poppy - 6-9-2012 at 19:46

You aware how many acid you must put to displace the reaction? haha

elementcollector1 - 6-9-2012 at 20:00

Not much: http://www.youtube.com/watch?v=zP9qEiaL4kQ

blogfast25 - 7-9-2012 at 05:16

Quote: Originally posted by elementcollector1

Not much: http://www.youtube.com/watch?v=zP9qEiaL4kQ

At a very minimum, check pH after acid addition. And if the chromate/dichromate concentration is relatively low the colour change may be subtle...

[Edited on 7-9-2012 by blogfast25]

elementcollector1 - 7-9-2012 at 08:45

It's not even subtle, still at that same bright yellow. Unless that's iron?
Here's an idea, why not run electrolysis of the chromate. That would get rid of any iron present as the insoluble hydroxide, and in all likelihood add more chromium to the solution from all that acid. Granted, I'll have to filter the mess again, but I'll boil it down this time to concentrate whatever chromate is present. Is there any way to remove trace iron that somehow made it into solution? And why didn't the concentrated HCl have any effect on the chromate?
I don't have anything in the way of pH checking.

blogfast25 - 8-9-2012 at 05:02

Quote: Originally posted by elementcollector1

It's not even subtle, still at that same bright yellow. Unless that's iron?

If it's iron it would have to be quite a lot to match the intense yellow of the CrO42- ion. Test for Fe (III) with KSCN (red FeSCN2-

and for chromate with lead nitrate (yellow PbCrO4 precipitates) or silver nitrate (reddish Ag2CrO4 precipitates).

But going by your description there should be no iron insolution because Fe(OH)3 is extremely insoluble.

Check pH of your chromate/dichromate solution.

You can also check for dichromate as follows:

* to a small sample add an excess of clear 'denaturated spirits' (ethanol + methanol), warm up: the yellow should dissapear and make way for a light green/blue (Cr3+)

* in the above substitute the alcohol with H2O2.

[Edited on 8-9-2012 by blogfast25]

elementcollector1 - 9-9-2012 at 11:54

Upon standing for three days, the solution now seems to have a metallic substance floating on top of it. I don't think this is metallic chromium, but what else could it be?

madcedar - 10-9-2012 at 06:46

Quote: Originally posted by elementcollector1

I read somewhere on this forum (thread was complete with pictures) that the decomposition of ammonium dichromate does not produce pure Cr2O3, it instead produces a rather dark green powder ripe with impurities
(Ze search function fails me yet again, I can't find ze thread...)

Plante1999 killed some spoons here:
http://www.sciencemadness.org/talk/viewthread.php?tid=15839

Look at these too:
http://www.sciencemadness.org/talk/viewthread.php?tid=19202
http://www.sciencemadness.org/talk/viewthread.php?tid=13347

elementcollector1 - 10-9-2012 at 08:52

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.

blogfast25 - 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]

elementcollector1 - 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.

blogfast25 - 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.

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

elementcollector1 - 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.)

blogfast25 - 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]

elementcollector1 - 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?

blogfast25 - 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]

elementcollector1 - 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...

blogfast25 - 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]

triplepoint - 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.

elementcollector1 - 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).

blogfast25 - 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]

elementcollector1 - 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.

blogfast25 - 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…

Poppy - 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.

elementcollector1 - 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).

blogfast25 - 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]

elementcollector1 - 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]

Poppy - 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.

elementcollector1 - 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?

blogfast25 - 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.

elementcollector1 - 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)
^Most useful online tool, ever.

blogfast25 - 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.

tetrahedron - 16-10-2012 at 18:02

Quote: Originally posted by elementcollector1

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 =)

elementcollector1 - 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...)

weiming1998 - 17-10-2012 at 19:13

Quote: Originally posted by elementcollector1

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...)

Using bicarbonate to directly precipitate Cr(OH)3 will work. Any potential Cr (III) carbonates/bicarbonates formed can be decomposed to the hydroxide easily by heating the solution after the bicarbonate had been poured into the acid and the solution stops fizzing.

elementcollector1 - 17-10-2012 at 19:36

Well, I have the bicarbonate/carbonate solution dripping through a sep-funnel into a good-sized flask overnight, we'll see what has happened come morning.
Incidentally, the reason I couldn't add the carbonate solution all at once was because the bubbling was so strong that I thought it might overbubble from the flask I have it in. I'm guessing the huge formation of bubbles is due to a bunch of CO2 from the reaction between the sulfuric acid (from the acidification + reduction step) and the sodium carbonate / bicarbonate (both basic). There's only a small amount of the Cr (III) sulfate solution in the flask, so I'll probably have to repeat this a couple times.

Incidentally, can anyone give me a picture of what fresh Cr(OH)3 looks like? Everyone describes it as 'grayish-green," or some variation, but I've never actually seen a picture. I always assumed it was just a grayer shade of the deep green of Cr2O3.

elementcollector1 - 25-10-2012 at 10:30

Well, I have no Cr(OH)3, and my solution (after quite a bit of dilution because of all the sodium carbonate/bicarbonate solution used) is now blue-green. I'm assuming this is chromite, so can I still prepare the hydroxide from this?
(I tried neutralizing the acid beforehand with NaOH, and all it did was immediately turn back to chromate.)
So, boil it down, and try again with more concentrated solutions?

blogfast25 - 25-10-2012 at 13:45

Quote: Originally posted by elementcollector1

I'm assuming this is chromite, so can I still prepare the hydroxide from this?

Cr(III) oxide (or hydroxide) is amphoteric: it can be dissolved in alkali (yielding chromites) or acids (yielding Cr(III) salts). Careful neutralisation of chromite will precipitate Cr(OH)3 but add too much acid and it will dissolve again...

elementcollector1 - 1-12-2012 at 17:50

I'm happy to report that I have a precipitate of light green Cr(III) sulfate from the reduced Cr(III) solution boiled down (there's still supernatant, though). I can try test tube-precipitations of this saturated solution with ammonia (speculated to work well, due to the poor amphoterism of Cr(OH)3 in ammonia which can be removed completely by boiling), sodium bicarbonate, and sodium carbonate.
This is pure chromium sulfate, correct? The acetone produced from reduction should have boiled off, as should the acid. My only concern is Na+ ions, which shouldn't be much of a problem considering I'm after an insoluble end-product.

elementcollector1 - 8-1-2013 at 21:20

Unfortunately, that precipitate was probably mostly sodium sulfate. Going to try concentrating it, removing as much sodium sulfate as possible, boiling to dryness to get rid of any remaining acid, and react with a few weak bases I have on hand to precipitate Cr(OH)3. Failing that, does anyone know a good method to convert sodium chromate or chromium sulfate to ammonium dichromate? I want to try the 'chemical volcano' anyways