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Author: Subject: Reactions of dichromates
Darkblade48
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[*] posted on 11-12-2005 at 14:37
Reactions of dichromates


Normally, I wouldn't play around with (hexavalent) dichromates due to their carcinogenity, but I was curious as to the various colours chromium (and other transition metals) have in different conditions.

Firstly, I added some potassium dichromate to a solution of sodium sulfite, and the solution went from a bright orange colour to a dull green colour (assuming the +6 to the +3 state, which would explain the bright green colour).

However, is the colour simply due to the Cr+3 ions in water? Or are they perhaps complexed with the water molecules to give this green colour?

Is there any possibility that I may have formed chromium (III) oxide (Cr2O3?)
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Mephisto
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[*] posted on 11-12-2005 at 15:36


You've produced Cr(III) as hydrate complex. I think the heat of the redox reaction might cause the formation of a green hydrate isomer. This isomerism of Cr(III)-hydrate complexes is known from chrome alum solutions, which are turning from violet to green when heated above 70 °C. It takes some days until the violet hydrate isomer is formed again. When using sulfuric acid and ethanol instead of the sodium sulfite for reducing Cr(VI) to Cr(III), you would get pure chrome alum (chromic potassium sulfate) after removing the aldehyde.

~Mephisto

Edit: No, I don't think there was any chromium(III) oxide formed.

[Edited on 11-12-2005 by Mephisto]




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Darkblade48
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[*] posted on 11-12-2005 at 18:57


Quote:
Originally posted by Mephisto
When using sulfuric acid and ethanol instead of the sodium sulfite for reducing Cr(VI) to Cr(III), you would get pure chrome alum (chromic potassium sulfate) after removing the aldehyde.


So in this case, using an acidified solution of ethanol would reduce the Cr(VI) to Cr(III), and I would get the pure chrome alum (which is supposed to be violet in colour?). Also you mention an aldehyde, would the OH group not be oxidized completely to the COOH functional group? I know that Jones reagent (CrO3) has the Cr in the +6 state, and this seems to oxidize the OH groups to COOH groups.

Quote:
Originally posted by Mephisto
Edit: No, I don't think there was any chromium(III) oxide formed.

On a totally different note, is it even remotely possible to get chromic oxide from the (potassium) dichromate, or is it only possible through an ammonium dichromate (thermal) decomposition?
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Mephisto
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[*] posted on 12-12-2005 at 07:47


Quote:
Originally posted by Darkblade48
Also you mention an aldehyde, would the OH group not be oxidized completely to the COOH functional group? I know that Jones reagent (CrO3) has the Cr in the +6 state, and this seems to oxidize the OH groups to COOH groups.


You're right. I've read the thing with the aldehyde in a book, in which it was the goal to produce chrome alum. Therefore the writer didn't care about the further oxidation of the acetaldehyde to acetic acid. The formed chrome alum (which crystallize from that solution) would be pure enough with only little traces of the by-products in its crystalline structure.

If you want do oxidise a primary alcohol just to the aldehyde, the Jones reagent isn't selective enough. Except for the case, in which the formed aldehyde is immediately distilled off from the reaction flask, to prevent a further oxidation.

Quote:
Originally posted by Darkblade48
On a totally different note, is it even remotely possible to get chromic oxide from the (potassium) dichromate, or is it only possible through an ammonium dichromate (thermal) decomposition?


Yes, potassium dichromate decomposes at 400 °C to chromium(III) oxide, potassium chromate and oxygen.

4 K2Cr2O7 → 2 Cr2O3 + 4 K2CrO4 + 3 O2

Additionally the Merck Index says, that chromium(III) oxide can be formed by the reaction of sodium dichromate or chromate with sulfur.

~Mephisto




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Darkblade48
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[*] posted on 12-12-2005 at 20:30


Ah, an interesting reaction I noted today. Mixing a solution of potassium dichromate with a solution of oxalic acid yielded (what I believe to be) a reduction of the Cr(VI) to a Cr(III) state, as the solution went from a bright orange to a dark, muddy brown colour.

However, I noticed the evolution of gas, would this be O2?
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[*] posted on 12-12-2005 at 22:43


With oxalic acid and potassium dichromate you get an tris(oxalato)chromate(III) complex. The evolved gas is CO2. With additional KOH you will get tripotassium trioxalatochromate(III), which is a nice substance for crystal growing. I tried this myself some years ago.

~Mephisto




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Darkblade48
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[*] posted on 13-12-2005 at 08:39


Most interesting. Is it possible to use a base other than KOH (i.e. NaOH) and get the respective sodium salt instead?

If not, how would I go about making the tripotassium trioxalatochromate (III) solution for crystal growing? Do I simply need to mix in the oxalic acid with the dichromate and wait for the reaction to occur before adding in the KOH? I assume then allowing the solution to stand and evaporate will allow crystals to form?
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[*] posted on 13-12-2005 at 11:22


Quote:
Originally posted by Darkblade48
Most interesting. Is it possible to use a base other than KOH (i.e. NaOH) and get the respective sodium salt instead?


So you would get a double salt like KNa2[Cr(C2O4)3]3 with 1 K+ from the potassium dichromate and 2 Na+ from NaOH. Okay, but this may influence the crystall shape in an unwanted manner. Maybe you are able to find simply potassium carbonate instead of KOH.

You can calculate the needed amounts from the following equitation:

4 KOH + K2Cr207 + 9 C2O4H2 · 2 H20 → 2 K3[Cr(C204)3]3 · 3 H2O + 6 CO2 + 23 H2O

The preparation is very simple. I've wrote down a detailed description in German language on www.LambdaSyn.com/synfiles/kaliumtrioxalatochromat.htm. Here is a quick translation:

284 g potassium carbonate or an equivalent quantity of potassium hydroxide are dissolved in 700-1000 ml water. 1165 g oxalic acid dihydrate are added slowly in small portions to the solution (CO2-formation ...). The solution is heated up to 80 °C. Under strong agitating 302 g potassium dichromate are added in smallest portions. The heating source is meanwhile switched off. It is important to add the potassium dichromate slowly and in small quantities, as the carbon dioxide evolvement can be very violent. This reaction is exothermic, too. After all potassium dichromate was added, the solution is boiled for one hour. It is hot filtered into container and left there for one week at room temperature. After this time large black crystals of tripotassium trioxalatochromate(III) have formed.

The fine spray, which escapes during the CO2-formation contains chromium. The preparation should be therefore done outside or in a fume hood.

~Mephisto




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[*] posted on 13-12-2005 at 19:45
There is still more


1. I read on a webpage about synthesizing of pigments that you can create a hydrated chromium oxide from potassium dichromate and boric acid. You melt them together at red heat for an hour or so then extract the potassium borate with boiling water.

2. A process simular to the oxalate complex is reduction of the dichromate with six moles of thiocyanate to form the crystalline Potassium Trithiocyanchromate.

3. You can neutralize with KOH to get the Chromate and add strontium chloride and ammonia to precipitate strontium chromate. Strontium Chromate is slightly soluble in boiling water so you can make a suspension of the strontium chromate in the equivalent amount of sodium bisulfate to get sodium dichromate and strontium sulfate.




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[*] posted on 13-2-2007 at 17:20


Could the procedure mentioned by Mephisto be done with the Na salts rather than K, so Na2Cr2O7, NaOH, and Oxalic Acid, to make Sodium Tris(Oxalato)Chromate (III)?



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[*] posted on 14-2-2007 at 03:33


Most likely it will work, but I doubt that you get nice crystalline results. Frequently, the sodium salts do not crystallize well and are very hygroscopic. So, isolation of the salt may be a real pain.



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[*] posted on 14-2-2007 at 16:23


Actually, I tried this myself last night, It formed crystals overnight, and for once, they are not hygroscopic!



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[*] posted on 15-2-2007 at 10:24


Are you sure these crystals are of the complex? I also did experiments with Na-salts and oxalato complexes, but I usually ended up with crystals of sodium oxalate, which were colored, due to impurities of the transition metals. Sodium oxalate is only sparingly soluble, IIRC 3 gram / 100 ml.



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[*] posted on 15-2-2007 at 22:49


This dissolves fairly well in water, I cannot say the exact solubility yet as I am waiting to let crystals finish drying. The crystals are just extremely dark, I do not think that they are Na oxalate. I will post a picture of them when I am finished however, and let you be the judge, as you seem to have done quite a lot with chromates and dichromates,



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[*] posted on 16-2-2007 at 00:21


Yes, if you have some pictures, that would be nice. And of course, I also can learn from this ;)
It is my experience that Na-salts usually are much more hygroscopic than K-salts, so if this is a nice exception, then I am always willing to adjust my ideas about this. I'll wait for your pics.




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