Panache
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Bright orange potassium ferrocycanide
Bought some of this last week as it was a discontinued line, however when i looked it up on wiki it is shown as bright yellow, the supply i obtained
was bright bright orange. Did web and site search no one ever talks about orange stuff, aqueous solutions are yellow however.
Wiki also states
'This compound is a strong reducing agent and is thus incompatible with oxidizing agents.[4] Addition of metal chlorates, perchlorates, nitrates, or
nitrites to a solution of carefully prepared and otherwise stable potassium ferrocyanide may result in a large explosion.[3]'
Is this true?
[Edited on 18-4-2010 by Panache]
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Ephoton
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I used to get that stuff ages ago.
it was orange as well. it sounds like what we were getting
was anhydrous and the yellow crystal is a hydrate.
I dont know about explosions though.
I used to put it with sulfuric acid and heat it.
it would go blue and release HCN 
[Edited on 18-4-2010 by Ephoton]
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woelen
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Are you sure you bought potassium ferrocyanide? The normal form in which it can be purchased is K4Fe(CN)6.3H2O and this is a pale
yellow crystalline solid. The anhydrous form is a nearly white (even more pale) powder.
You might instead have K3Fe(CN)6. This compound is a crystalline orange/brown solid and it gives bright yellow solutions in water.
http://woelen.homescience.net/science/chem/compounds/potassi...
http://woelen.homescience.net/science/chem/compounds/potassi...
Both compounds give HCN when heated with sulphuric acid of moderately high concentration. Simply adding the compounds to the acid without heating does
not release HCN.
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Ephoton
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thanx for that info.
I honestly thought what I had was ferro.
I got it from vanbar and thats how they had it advertised
at the time.
I would use them unless you pick the stuff up your self as they have ripped me off again not sending product and not replying to my emails.
they were happy to take my money though.
[Edited on 18-4-2010 by Ephoton]
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S.C. Wack
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Both are color-sensitive to all manner of contamination.
The ferricyanide is more red than orange. The anhydrous ferro- is white, and recrystallizes from water in large pale crystals, which have no color
other than yellow.
The calculated amount of aq. KOH and excess H2O2 reduces ferricyanide quantitatively, acid and H2O2 gives large red crystals from the ferrocyanide.
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The WiZard is In
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Quote: Originally posted by Panache  |
'This compound is a strong reducing agent and is thus incompatible with oxidizing agents.[4] Addition of metal chlorates, perchlorates, nitrates, or
nitrites to a solution of carefully prepared and otherwise stable potassium ferrocyanide may result in a large explosion.[3]'
Is this true?
[Edited on 18-4-2010 by Panache] |
Explosion? Yup. In years long gone by K ferrocyanide w/
potassium chlorate &c., was suggested for use in a
number of explosives. All of which were v/ v/ v/ sensitive
and not usable.
Instructions for compounding one — advised mixing the
ingredients with a feather! [Ref available on request.]
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Rosco Bodine
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| Quote: Originally posted by Panache | Bought some of this last week as it was a discontinued line, however when i looked it up on wiki it is shown as bright yellow, the supply i obtained
was bright bright orange. Did web and site search no one ever talks about orange stuff, aqueous solutions are yellow however.
Wiki also states
'This compound is a strong reducing agent and is thus incompatible with oxidizing agents.[4] Addition of metal chlorates, perchlorates, nitrates, or
nitrites to a solution of carefully prepared and otherwise stable potassium ferrocyanide may result in a large explosion.[3]'
Is this true?
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No, at least not precisely true anyway. What may more likely happen is an oxidation of the ferrocyanide in solution to the ferricyanide in solution.
Both the dry ferrocyanide and the dry ferricyanide may function as fuels and form sensitive, pehaps unpredictably spontaneously deflagrating, easily
ignitable, or detonating mixtures with potassium chlorate in particular, and probably also reactive with other oxidizers to a probably lesser extent.
Sometimes 30 minutes in the lab is better than 30 hours in the library. But what is deeper than the heart of Texas is six inches into Virginia. 
http://www.youtube.com/watch?v=gjykMVzFd7g&fmt=18 Coventry Carol
[Edited on 19-4-2010 by Rosco Bodine]
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Bolt
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Oddly sexual.. but I don't get it.
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Rosco Bodine
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a contrast of dry and wet humor ....something like the differing conditions which produce explosions or not,
but that was a touche' meant for the wiz
and the song ....well it was just a song to salve the wound
All chemists should vigorously deny that there is ever any seductive behavior on the part of any researchers, students, or staff in laboratories,
offices, or supply rooms ...as that would be an unacceptable breach of scientific decorum,
whilst all exaggerations possible of the fatal dangers which should be avoided by the public at such facilities cannot be overemphasized in the
interest of maintaining an acceptable level of privacy for the proper conduct of scientific pursuits
with sufficient privacy..... and uninterrupted till those activities satisfactory completion.
[Edited on 19-4-2010 by Rosco Bodine]
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The WiZard is In
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| Quote: |
Oddly sexual.. but I don't get it. |
Hint — What is the Latin word for a knife sheaf?
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Rosco Bodine
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It's enough to get a rise out of any man worth his salt, be it ous or ic
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Panache
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The crystals i have match your photograph quite evenly woelen, i guess i better go and research this compound a little.
Ta.
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kmno4
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K4Fe(CN)6 x 3H2O can be "bright bright orange" especially if it is freshly prepared and in form of not too small crystals.
If partially dehydrated and/or in a form of fine powder it is clearly yellow and if anhydrous - white.
Of course chemical test with any soluble Fe(III) salt is better than colour test.
But you can also warm small amount in 90-100 C and see if it looses water and becomes white.
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woelen
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I do not agree with kmno4. I have seen crystals of K4Fe(CN)6.3H2O of more than 1 cm diameter and these crystals were pale yellow and shiny. Nothing
orange at all. If the ferrocyanide is orange, then it is impure, the pure salt is pale yellow (or white if anhydrous).
The ferricyanide also can form large crystals. These crystals are beautifully deep red/orange with a brown hue.
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Panache
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well a solution (yellow) remained as such after addition of FeCl3, however it also remained yellow after i dropped in a very rusty razor blade, i know
the oxides of Fe(II) are very insoluble but surely one would expect Prussian blue to evolve to some extent. Perhaps one must wait for the [Fe(ii)] to
negate the Fe(III) i have added previously this could take some time.
A quick note the CRC and wiki do mention any decomposition of the salt at its melting point (300C), can others corroborate this? Also Woelen on your
site you mention ferrocyanides instability to UV, do you have more information on this, what are the photolysis products and what wavelengths are
involved?
In the time this note has taken to write i believe the yellow has begun to adopt the slightest green hue, howevery i'm likely only looking for it.
Wiki makes no reference to any acidic forms of the ferricyanide/ferrocyanide anion, do they exist, perhaps at low temperatures, as in directly
protonated by H+(lewis bronstead if i remember correctly), i understand as the potassium salt itself acts as a weak Lewis acid affecting organic
reactions catalytically in this regard.
The razorblade now has a distinct blue line running along it edge, very sharp colour evolution. Lovely!
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entropy51
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| Quote: Originally posted by kmno4 | | K4Fe(CN)6 x 3H2O can be "bright bright orange" especially if it is freshly prepared and in form of not too small crystals. |
Only if they are really ferricyanide, as any chemist knows. Such mistakes in labeling are common in certain countries in eastern
Europe, and account for some of the ridiculous "data" reported from that region.
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kmno4
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"well a solution (yellow) remained as such after addition of FeCl3, however it also remained yellow after i dropped in a very rusty razor blade, i
know the oxides of Fe(II) are very insoluble but surely one would expect Prussian blue to evolve to some extent."
Dilute solution of K3Fe(CN)6 (5%) is very sensitive to Fe(II) but you have to add small amount of HCl(aq) to this solution. 1 drop per 10 cm3 is good.
Then even stirring this sol. with a Fe nail, rasor blade, spoon (SS), knife... etc. gives blue coloration.
ps.
| Quote: | | Only if they are really ferricyanide, as any chemist knows. |
"Never argue with an idiot. The idiot will drag you down to his level and then beat you with his experience."
And I never will...
[Edited on 21-4-2010 by kmno4]
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woelen
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Panache, your observation is another indication that you have the ferricyanide and not the ferrocyanide. Ferricyanide contains iron in oxidation state
+3 and when a ferric salt is added to that, such as FeCl3 then a weak complex is formed, which has a somewhat brown color. In your strong yellow
solution you do not really see that. At higher concentrations you may notice it as a brown coloration.
The deep blue color only is formed when either a ferrous salt is mixed with the ferricyanide or when a ferric salt is mixed with the ferrocyanide. You
need a mix of iron(II) and iron(III).
So, if you want to see the deep blue color, then dissolve some iron (e.g. from a nail) in hydrochloric acid. That solution certainly contains iron(II)
besides iron(III) and if you add a few drops of the ferricyanide solution, then you'll get a deep blue solution. No need to fear formation of HCN in
this reaction, the cyanides are sufficiently strongly coordinated to the iron that they are not released by cold hydrochloric acid.
Ferrocyanide gives a pale (almost white) precipitate with a pure iron(II) salt. This pale precipiate slowly turns blue, due to aerial oxidation.
Summarizing:
iron(II) + ferrocyanide --> pale off white precipitate, which easily is oxidized by oxygen from air
iron(II) + ferricyanide --> deep blue precipitate
iron(III) + ferrocyanide --> deep blue precipitate
iron(III) + ferricyanide --> brown complex formation, which remains in solution
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zed
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Explosions? Yes!
Years ago, a buddy demonstrated to me, an explosive powder produced using Potassium FerroCyanide as one component.
Detonates abruptly without any containment whatsoever. Bamm!
It seems unlikely however, that a dilute aqueous solution could actually detonate.
The water would probably "sink" away too much energy. Might release enough heat and gas, to burst a sealed glass vessel though.
[Edited on 27-4-2010 by zed]
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Formatik
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| Quote: Originally posted by Panache | Wiki also states
'This compound is a strong reducing agent and is thus incompatible with oxidizing agents.[4] Addition of metal chlorates, perchlorates, nitrates, or
nitrites to a solution of carefully prepared and otherwise stable potassium ferrocyanide may result in a large explosion.[3]' |
This is a good case in point demonstrating weighing statements on wikipedia with common sense and reference proofing. The original cited reference in
the wiki article (JTbaker) doesn't even mention anything about solutions.
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woelen
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The remark that ferrocyanide or ferricyanide is a dangerous compound with explosive properties is the same as telling that water is a dangerous
compound which may cause fire or explosion.
Water can cause fire or explosion when brought in contact with K-metal. Ferricyanides and ferrocyanides can cause fire or explosion when mixed with
chlorates (not with perchlorates!) or strong nitric acid. I even can make paper explode (I once did) by mixing it with Mn2O7.
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densest
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AFAIK, potassium ferrocyanide & potassium perchlorate will deflagrate pretty vigorously if ignited. This patent talks about it: http://www.freepatentsonline.com/3793100.html The mixture is pretty stable if not provoked.
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