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DJF90
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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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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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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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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.
Chemistry is life (and a whole lot more)
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Jor
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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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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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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! :-)
"Don't run away....I come in peace!!"
[Edited on 4-6-2009 by Elawr]
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woelen
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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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Nice experimentation woelen!
Chemistry is life (and a whole lot more)
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pHzero
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| 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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woelen
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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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