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Author: Subject: Ferricyanide precipitation?
DJF90
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[*] posted on 2-6-2009 at 17:22


What I think may be possible is that an electron in the ground state can be promoted to an excited state by the influx of light, and then an electron can be accepted into that lower energy orbital (the one the excited electron was in), thus reducing the metal species. This would probably be a sensible theory for the [Fe(OH2)6](3+) species.

However I dont think this possible with the ferricyanide complex - the cyano ligand is high in the spectroscopic series meaning that the crystal field splitting parameter, delta oct., is large (i.e. d5 low spin), effectively prohibiting the excitation of an electron in the t2g set. However accepting one electon will fill the t2g set (d6 low spin), producing a gain in ligand field stabilisation energy (by 0.4 delta oct.; quite a substantial gain providing that delta oct. is very big due to the six cyano ligands). I suspect it is this gain in LFSE that makes the reduction of ferricyanide to ferrocyanide so easy.

[Edited on 3-6-2009 by DJF90]
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Elawr
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[*] posted on 2-6-2009 at 21:35


I agree with Woelen in that we should not be ugly to phZero, particularly in the manner of DJF90. This is clearly a bright and curious young man who is starting out the same way many of us did, long ago when we were young.

How many of us learned the hard way about using our thumbs as stoppers? And, does any one not agree that it is surely a small miracle that many of us (self included) survived our early days of chemical adventures without permanent injury from poisoning, or hideous disfigurement from fire or explosion?

By the way...I think phzero handled himself quite admirably as far as DJF90's flame attack is concerned.

I suggest we not be so hard on this young man, so that we don't run him off. He is capable of learning.

Besides, one day we might be working for him :-)




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DJF90
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[*] posted on 2-6-2009 at 23:27


I have nothing against pHzero, so long as he is responsible and does not take unnecessary risks. You forget that when each one of us experiments, should something go badly wrong then this puts chemistry as a hobby at serious risk for the rest of us. I would just appreciate a little common sense from this "bright and curious young man", as I dont want my priviledges (i.e. being able to source chemicals and experiment) taken away from me because of his ignorance. He is not the only one, granted, but it would be nice if he looked after himself for the sake of all of us. I'm sorry if you feel like I've been flaming you pHzero.

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Hydragyrum
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[*] posted on 2-6-2009 at 23:45


Quote: Originally posted by DJF90  
However I dont think this possible with the ferricyanide complex - the cyano ligand is high in the spectroscopic series meaning that the crystal field splitting parameter, delta oct., is large (i.e. d5 low spin), effectively prohibiting the excitation of an electron in the t2g set. However accepting one electon will fill the t2g set (d6 low spin), producing a gain in ligand field stabilisation energy (by 0.4 delta oct.; quite a substantial gain providing that delta oct. is very big due to the six cyano ligands). I suspect it is this gain in LFSE that makes the reduction of ferricyanide to ferrocyanide so easy.

What you say makes sense; I had forgotten about the effect of cyanide, but I believe you are correct and that the complex is always low-spin. The only problem remaining then is why would it need light to 'activate' for reduction? (this from previous comments)

Perhaps, as you say, light is not needed after all.




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Jor
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[*] posted on 3-6-2009 at 00:40


I agree Elawr, but how did we learn to experiment safely? Probably after mistakes or tips from other more experienced chemists. After this single flame comment he probably won't do it again. So he learned right? That how we learned it all as well probably, because someone tells you, possibly in a harsh way, to not experiment like that.
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woelen
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[*] posted on 4-6-2009 at 12:13


I have had the mix of solutions of KOH and K3Fe(CN)6 brewing now for more than one day and still the solution is as clear as when I prepared it. No precipitate at all. I, however, had the liquid in a dark place. Now I let it stand in a brightly lit room. Unfortunately the next few days will be very cloudy and dark, so I might need to add some artifical light (black light, contains a lot of near UV). Within a day or two I will let you know how the solution behaves in light.



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Elawr
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[*] posted on 4-6-2009 at 15:29


Good, valid points both of you DJF90, Jor...

I would hate to see the boy get run off from here; that's all.

A good chemist never stops learning and uses all possible sources,... and sometimes a little flaming is a good thing, when applied judiciously! :-)








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[Edited on 4-6-2009 by Elawr]




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woelen
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[*] posted on 6-6-2009 at 08:36


I had the solution mix in sunshine now for a few hours and it really is amazing to see how quickly it becomes turbid and a precipitate is formed. Within 1 hour, the solution becomes turbid and some material sticks to the glass. After four hours of bright sunshine, the solution is turbid and orange/brown and a thin layer of orange/brown solid material has settled at the bottom.

So, the reaction indeed requires light. In the dark no decomposition occurs, but in the light it goes fairly quickly.

I also added some of the yellow liquid with orange/brown precipitate to dilute reagent grade hydrochloric acid. When this is done, then the liquid becomes clear again, and the solution becomes green, but no intense blue coloration can be observed. So, there only is a very small amount of iron(II). I think that the light makes conversion of iron(III) to iron(II) possible, accompanied with partial breakdown of the cyanide complex but that subsequently the iron(II) is oxidized back to iron(III), possibly by oxygen from the air, or by some species in solution. This results in formation of Fe(OH)3 as solid matter and explains why there is so little iron(II) in solution.




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Hydragyrum
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[*] posted on 6-6-2009 at 14:50


Nice experimentation woelen!



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pHzero
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[*] posted on 6-6-2009 at 15:09


Quote: Originally posted by woelen  
I had the solution mix in sunshine now for a few hours and it really is amazing to see how quickly it becomes turbid and a precipitate is formed. Within 1 hour, the solution becomes turbid and some material sticks to the glass. After four hours of bright sunshine, the solution is turbid and orange/brown and a thin layer of orange/brown solid material has settled at the bottom.

So, the reaction indeed requires light. In the dark no decomposition occurs, but in the light it goes fairly quickly.

I also added some of the yellow liquid with orange/brown precipitate to dilute reagent grade hydrochloric acid. When this is done, then the liquid becomes clear again, and the solution becomes green, but no intense blue coloration can be observed. So, there only is a very small amount of iron(II). I think that the light makes conversion of iron(III) to iron(II) possible, accompanied with partial breakdown of the cyanide complex but that subsequently the iron(II) is oxidized back to iron(III), possibly by oxygen from the air, or by some species in solution. This results in formation of Fe(OH)3 as solid matter and explains why there is so little iron(II) in solution.


Ah, interesting findings. Seems a lot more likely than my suggestion. So if the ferricyanide was broken down into iron and cyanide, perhaps this could a a new, safer way of getting cyanide ions?
The normal way to get CN- in the lab is to acidify ferricyanide and bubble the HCN through and alkali, right? If you added some KOH to ferricyanide and kept filtering it, the Fe(OH)3 would get stuck in the filter, pushing the equilibrium to the KCN+Fe(OH)3 side. That way you dont have to handle CN- in the gas phase

Well, all of this is hypothetical. I'm certainly not trying that, I'm just trying to find possible uses for this.

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[*] posted on 7-6-2009 at 06:34


I do not think that this has any practical uses. The reaction is slow and incomplete and I'm not sure that cyanide is one of the reaction products (I am inclined to think that cyanogen is formed, which immediately is converted to cyanide and cyanate, which in turn hydrolizes to ammonia and carbonate). This reaction, however, is interesting from a theoretical point of view and that's why I tried it with different conditions.



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