Ferricyanide precipitation?

No, you didn't get it

KOH does not reduce iron(III) to iron(II), it is not a reductor at all. Ferrocyanide also is soluble to a large extent and this does not precipitate from the solution. You did not get K4Fe(CN)6. K4Fe(CN)6 is light yellow/white and not orange.

I can imagine that in very strong alkaline solutions some of the ferricyanide does decompose, albeit only a little bit. The anti-caking agent in kitchen salt is ferrocyanide, not ferricyanide. The ferrocyanide ion indeed is very stable and the cyanides cannot be decoupled from the iron. Ferricyanide is less stable. It does not easily loose cyanide, but I can imagine that it did in your experiment.

The equation you gave in your first post can be regarded as an equilibrium, which is far to the left, but not 100%.

Quote: Originally posted by woelen

I can imagine that in very strong alkaline solutions some of the ferricyanide does decompose, albeit only a little bit. The anti-caking agent in kitchen salt is ferrocyanide, not ferricyanide. The ferrocyanide ion indeed is very stable and the cyanides cannot be decoupled from the iron. Ferricyanide is less stable. It does not easily loose cyanide, but I can imagine that it did in your experiment.

Quote: Originally posted by entropy51

pHzero, one doesn't ordinarily think of NaOH as a reducing agent, but in this case you could just be correct. On page 223 of the Textbook of Inorganic Chemistry (in the Forum library), Volume IX Part II by Friend, page 223, it's stated that potassium ferricyanide is reduced to potassium ferrocyanide in alkaline solution. On page 135 of The Chemistry of the Cyanogen Compounds by Williams ( available at books.google.com) it is stated that the ferricyanides are reduced to ferrocyanides by caustic alkali. I've been meaning to try this myself, but you beat me to it. Maybe you could try some of the tests for ferrocyanide (perhaps formation of Prussian Blue) and see if you can verify the references.

Quote: Originally posted by entropy51

Well something happened, but what? Try making a solution of ferricyanide and adding equal amounts to two test tubes. Add a solution of KOH to one tube and an equal amount of water to the other. Let the reaction stand a few minutes and add equal amounts of a ferric solution to both tubes. That way you can compare the two. You didn't use your thumb as a stopper when you shook that tube with KOH in it, did you? That's not a good idea. Use a cork or a rubber stopper.

Yep I did, its all I've got unfortunately - no stoppers
It's ok though - I've shaken saturated solutions of KOH with my finger on the end and the only problem is it stings a little if you get it it in a cut.

I'll try that on monday probably. I know that was a very bad experiment - no control and I mixed the reactants as solids.

Quote: Originally posted by DJF90

Are you being stupid on purpose or is this just a general trait of your personality? You suggest that KCN is made, yet you shake the aqueous solution that could contain CN-. and also OH-, using your thumb as a stopper. It is just bad practice, no mater what chemicals are involved. You don't seem to experiment in a safe way and I urge you to reconsider your methodology before you end up seriously injuring yourself or those around you. And I sincerely hope you was wearing safety specs; potassium hydroxide solution of that concentration will readily blind you should your thumb slip off the opening of the test tube... KOH solution is pretty slippy after all. And I think theres probably a good reason why your parents get annoyed when you have chemicals in the house. [Edited on 31-5-2009 by DJF90]

Quote: Originally posted by DJF90

Then you are more idiotic than I thought; To have safety specs and not wearing them, especially when using chemicals like conc. KOH solution that can, quite rapidly (we're talking <10-15 seconds here) blind you, is just plain irresponsible. Surely you can understand why your parents are pissed; imagine what the situation would be if you had dropped something that has a more hazardous effect, rather than merely cosmetic (perhaps something like mercury metal... now that would have been a disaster). And its common sense to use smelly chemicals outside if you dont have a fume hood. Chemistry is a serious subject; getting it wrong has very serious and very real consequences, and it'll do you the world of good to perhaps consider this before experimenting further.

Quote: Originally posted by DJF90

Finally someone can see exactly what I'm saying! Its idiots like pHzero that got decent chemistry sets removed from the shops. Now they're full of crap "chemicals" and not even glassware last I heard... plastic test tubes.. fancy that!

Should we really bash a starting newcomer this way? I agree with all the people, who tell that you must take safety measures and if you use KOH of high concentration then it is wise to use goggles. But now everyone is falling over the other screaming how irresponsible pHzero has been and that is not good at all for home chemistry neither. Who of us has not done unwise things when you were 15 or so ? So, one remark is enough and then things are said. In this way we are not better than those "politically correct" people who are despised by so many of us.

So, the point is made now and let's go on about the chemistry of this interesting subject. Because of the stoppering with the thumb, I can imagine that material from the skin served as reductor and allowed some of the ferricyanide to be converted to ferrocyanide.

So, this experiment must be repeated in a clean test tube, with both chemicals dissolved separately with careful swirling (and no shaking) and then mixing the two solutions and keeping them in a dark place.

Btw, potassium ferricyanide decomposes in aqueous solution anyway, even a plain water solution decomposes, especially in the presence of light. Bright sunlight causes decomposition in minutes.

The mechanism behind this is that Fe(III) easily is activated by light and then can give off an electron. This is the basis for iron-based photography (cyanotype, old noble process). Just for fun, mix some ferricyanide, some iron(III) salt (without base added) and keep in the light. The solution turns blue soon. The cyanotype process exploits this by using a solution of ferric oxalate or ferric citrate and mix this with a solution of potassium ferricyanide and absorbing this in paper. Parts exposed to light become blue, other parts are not colored. Washing in very dilute acid (almost just water) then removes the unreacted iron(III) salts while the insoluble prussian blue sticks to the fibers of the paper.

For the above reason, it is best to store ferricyanides in a dark place and only as a solid. Ferrocyanides are much more stable and can also be kept in a light place. Iron(II) salts, like ferrosulfate even can best be stored in a bright place, preferrably in bright sunlight. This increases shelf life considerably, because any oxidized iron (to +3) is more easily reverted to its +2 state by the light.

Quote: Originally posted by woelen

The mechanism behind this is that Fe(III) easily is activated by light and then can give off an electron.

Sorry for the mistake. That was one of my unclear moments

I meant accepting an electron. Iron(III) is acitivated by light and then easily accepts an electron, going to iron(II).

Quote: Originally posted by woelen

Quote: Originally posted by DJF90   Finally someone can see exactly what I'm saying! Its idiots like pHzero that got decent chemistry sets removed from the shops. Now they're full of crap "chemicals" and not even glassware last I heard... plastic test tubes.. fancy that! Should we really bash a starting newcomer this way? I agree with all the people, who tell that you must take safety measures and if you use KOH of high concentration then it is wise to use goggles. But now everyone is falling over the other screaming how irresponsible pHzero has been and that is not good at all for home chemistry neither. Who of us has not done unwise things when you were 15 or so ? So, one remark is enough and then things are said. In this way we are not better than those "politically correct" people who are despised by so many of us. So, the point is made now and let's go on about the chemistry of this interesting subject. Because of the stoppering with the thumb, I can imagine that material from the skin served as reductor and allowed some of the ferricyanide to be converted to ferrocyanide. So, this experiment must be repeated in a clean test tube, with both chemicals dissolved separately with careful swirling (and no shaking) and then mixing the two solutions and keeping them in a dark place. Btw, potassium ferricyanide decomposes in aqueous solution anyway, even a plain water solution decomposes, especially in the presence of light. Bright sunlight causes decomposition in minutes. The mechanism behind this is that Fe(III) easily is activated by light and then can give off an electron. This is the basis for iron-based photography (cyanotype, old noble process). Just for fun, mix some ferricyanide, some iron(III) salt (without base added) and keep in the light. The solution turns blue soon. The cyanotype process exploits this by using a solution of ferric oxalate or ferric citrate and mix this with a solution of potassium ferricyanide and absorbing this in paper. Parts exposed to light become blue, other parts are not colored. Washing in very dilute acid (almost just water) then removes the unreacted iron(III) salts while the insoluble prussian blue sticks to the fibers of the paper. For the above reason, it is best to store ferricyanides in a dark place and only as a solid. Ferrocyanides are much more stable and can also be kept in a light place. Iron(II) salts, like ferrosulfate even can best be stored in a bright place, preferrably in bright sunlight. This increases shelf life considerably, because any oxidized iron (to +3) is more easily reverted to its +2 state by the light.

Firstly, thank you for your understanding that we all make mistakes, especially when we're young, and learn from them.
As I said, I've learnt from the experience and now I know to wear goggles etc. (At school, we had to wear goggles for an osmosis experiment where we put pieces of potato in sugar solution, so I've always assumed that theyre just for the sake of complying with H&S laws)

I'll try what you recommended tonight. It's a really sunny day and there isn't really anywhere dark to do it at the moment.

As for the cyanotype process, I've tried that I mixed FeCl3 (got loads of it cause I buy it in bulk and sell it on ebay as a PCB developer to make a bit of money) and ferricyanide then soaked a piece of paper in it. Then I put a few leaves on it and left it out in the sun for 5 minutes and I got a vaguely recognisable image.

Edit: actually I'll leave it till friday, when I'll have some rubber stoppers. And my mums bringing me some latex gloves home from work.

[Edited on 2-6-2009 by pHzero]

Quote: Originally posted by woelen

Sorry for the mistake. That was one of my unclear moments I meant accepting an electron. Iron(III) is acitivated by light and then easily accepts an electron, going to iron(II).

I'm not really good at the electron orbits, but what would come out of that energy given off by going to a lower spin. I'd assume that the blueness would come from the complex forming in the solution, which would change the Iron. Tell me if I'm wrong. I don't really understand all the spinning of the electrons (I've gone so far as to look at bonding and various organic bonds, but not really into the physical spin, or the physics part of it).

I understand . I've learned a lot from here.

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.

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.