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Author: Subject: Did I just capture some Cr(V)? [edit: Nope, but interesting discussion anyway.]
j_sum1
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[*] posted on 18-8-2021 at 14:16
Did I just capture some Cr(V)? [edit: Nope, but interesting discussion anyway.]


Mild curiosity here.

I was demonstrating for some students the oxidation of alcohols and differentiation between primary/secondary and tertiary. To complete the demo I showed that ketones and carboxylic acids do not oxidise. All experiments were done at test tube scale using Na2Cr2O7 / H2SO4.

When cleaning up I wanted to convert all Cr(VI) to Cr3+ as is my normal practice. To do this I added some ethanol to the test tubes and left them in the warm water bath.

Coming back later I noticed that my test with GAA was different from the others. Obviously it had esterified. But it had also turned a lovely pale purple shade. Cr(V) is the only thing I can think of that would do that, but it is notoriously unstable, especially if there is water present (as would be the case this instance.) The solution is holding its colour hours later. Any insights as to what it might be?

[Edited on 19-8-2021 by j_sum1]
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[*] posted on 18-8-2021 at 14:21


Depending on its ligands, Cr(III) can be purple or green. For example, chromium alum, KCr(SO4)2 is deep purple. A good way to make chrome alum is by reduction of K2Cr2O7 with ethanol in the presence of sulfuric acid.
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[*] posted on 18-8-2021 at 14:36


Makes sense. I forgot about chrome alum. And yes, the colour does match.

I just don't associate Cr with purples. Maybe I need to change.

What particular structures are associated with green / purple for Cr3+?
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[*] posted on 18-8-2021 at 16:28


Speaking of chrome alum, all you have to do to produce it from your waste solutions is evaporate them down! You have sulfate and equal numbers of chromium(III) and potassium ions dissolved, and chrome alum is by far the least soluble product.

I think I got this from woelen, and I'll try not to mess up the explanation: chromium(III) in its free state or with certain ligands on it is the typical violet color, but in solution it can trade those ligands for water molecules, forming a series of dark green aquo complexes. These are even more favored when certain counterions are present (chloride being one of them) and universally at elevated temperatures, which is why if you heat solutions of chromium(III) above about 40C they'll turn dark green and only very slowly, over a period of days or even weeks, revert back to blue/violet.

[Edited on 8-19-2021 by Amos]
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[*] posted on 18-8-2021 at 17:52


Ok. Except that I used sodium dichromate and not potassium dichromate. And the sulfuric acid is in very low concentration.

But basically you are saying that the relative absence of water causes the purple colour to be visible rather than the deep green/grey that I am familiar with.
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[*] posted on 19-8-2021 at 07:06


It is the other way around. Chromium(III) with only water ligands is purple (albeit somewhat vaguely purple, it also is somewhat bluish/greyish). If chloride ligands or sulfate ligands are attached, then it turns green.

A simple but nice experiment is to dissolve some chromium(III) sulfate or chrome alum in water, which gives you a purple/bluish solution. Then heat the solution, till near boiling. The solution turns green. Next, let it stand. On cooling down it remains green, but after several days the colors shifts again and after a week or so, the solution is purple/bluish again.

Acetate ligands give a deep purple color and even better results are obtained with oxalate. Dissolve a little K2Cr2O7 or Na2Cr2O7 in water and add an excess amount of oxalic acid (no other acid should be added, so that only oxalate ions can be coordinates besides the water). Gently heat. The color slowly turns from orange through shades of brown to intense purple. The latter is the color of the oxalato complex, formed in this reaction.




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