TmNhRhMgBrSe
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Is sulphite strong enough to reduce Cr(VI) to Cr(II)?
I mixed some chromate, acetic acid and sulphite. The solution turned blue. So is the chromium +2?
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DraconicAcid
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No, it's Cr(III).
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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TmNhRhMgBrSe
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Then why it's blue?
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Metallus
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This may help https://www.chemguide.co.uk/inorganic/transition/chromium.ht...
Cr3+ is green when some of the water molecules of the hexaquaion are replaced by sulphate/chloride ions, otherwise it's blue/grayish. Since you are
working in CH3COOH medium, you probably have mostly the pure hexaquaion. Try to add a good amount of a chloride or sulphate and see if the solution
stays blue or turns green.
Take into account that normally dichromates are used in conjunction with H2SO4 >2M which provides a good amount of sulphate ions, therefore giving
the green solution after reduction to Cr3+.
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teodor
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The discussion in this thread might also help.
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AJKOER
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As usual my fellow 'experts' could be myopic chemists.
Is it possible there is any presence and continuing O2 exposure or lab light? Smell of SO2? Was the sulphite stored in the dark from free oxygen
exposure?
Any dissolved oxygen in the presence of solvated electrons from the redox reaction, speculatively could introduce the superoxide radical anion. The
latter may further react as follows:
Cr(lll) + O2•- = Cr(ll) + O2 (the reaction is reversible)
Also, any SO2 from the sulphite, could further engage in the following equilibrium reaction:
SO2 + •O2- <--> •SO2- + O2 (see Eq (48) at https://pubs.acs.org/doi/full/10.1021/acs.chemrev.5b00407 )
where the formed sulfur dioxide radical anion is well noted for its reducing properties.
Light can also have an effect on the system and even pre-experiment impacted the chemical composition of the sulphite exposed to air (see https://pubs.acs.org/doi/pdfplus/10.1021/j150261a011 ).
So, before being dogmatic, I would suggest repeating the experiment after at least removing oxygen and lab light from the system, as well as checking
the sulphite, to confirm the claimed explanations.
[Edited on 30-7-2019 by AJKOER]
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woelen
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@AJKOER: Don't make things more complicated than they are. The presence of oxygen or light is totally irrelevant in this context. When the sulfite is
stored in a well sealed container, then it can be kept around indefinitely (I can say from experience, I have a 20+ year old stock of Na2SO3 and this
is as good as when I purchased it, i stored it in a thick-walled plastic jar).
What happens in this reaction is formation of complexes. Chromium has the specific property, that when it goes from oxidation state +6 to oxidation
state +3 (or from oxidation state +2 to oxidation state +3) that as soon as it reaches the +3 oxidation state, it picks its ligands and the resulting
complexes are quite inert.
Chromium forms green/blue sulfato complexes with sulfate ion, which are quite stable. As soon as the sulfite ion is ozidized to sulfate ion by the
chromate (better: dichromate in acidic solution), the resulting chromium(III) picks up sulfate ions and forms a sulfato complex.
In the absence of sulfate ions, chromium can form other colored complexes on transition to oxidation state +3.
In the persence of chloride ions, a chloro-complex is formed, which is purely green, it has no bluish hue.
In the presence of oxalate ions, an oxalato-complex is formed, which has an intense very dark purple/grey color.
In the presence of only non-coordinating ions, a fairly dark, somewhat indistinct grey/blue/violet color is obtained, the color of the aqua complex of
chromium(III).
If you want to see the formation of chromium complexes in action, but without redox reaction, do the following:
Dissolve some chrome alum or chromium(III) sulfate in water. The crystalline solid is dark purple/black, the solution has a somewhat vague
grey/blue/violet color, the color of the hexaqua complex of chromium(III). If you heat the solution to near-boiling, then it remains clear, but it
becomes beautifully bright green with a bluish hue. This is due to ligand exchange, water molecules are replaced by sulfate ions in the coordination
sphere of the chromium(III) ions. If you allow the solution to cool down again, then it remains green. It takes days, or even weeks if the room is
cold, before the original vague color of the hexaqua complex of chromium(III) appears again.
If you do the same experiment with chromium(III) nitrate, then you do not get a green color on heating. There is a slight shift in color, but this
only is due to a slight change of the aqua complex and on cooling down, the original vague grey/blue/violet color is obtained again. If some NaCl or
HCl is added to the solution and the experiment is repeated, then you get a pure green solution on heating. This is due to formation of a chloro
complex of chromium. As with the sulfato complex, this green color is quite stable and takes a long time to revert to the original color again.
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andy1988
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By coincidence I read this today:
| Quote: | | Chromous (chromium II), a biologically inactive valence state, exists, but once exposed to air it rapidly oxidizes to chromium III, a biologically
active form of chromium [1]. |
[1] Wexler, P., & Gad, S. C. (1998). Encyclopedia of toxicology. Academic Press. Volume 1 page 340
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