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pHzero
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[*] posted on 28-5-2009 at 15:51
Ferricyanide precipitation?


I just added KOH to a solution of K3Fe(CN)6. The solution didn't change colour, but a yellow-orange precipitate was formed.

I initally thought that this precipitation reaction had taken place:
K3Fe(CN)6 + 3KOH --> 6KCN + Fe(OH)3 (s)

But I'd also read that the Fe(CN)6 3- ion's very stable and cant be broken easily, and that ferricyanide's fairly non-toxic to the extent that it's safely added to salt as an anti-caking agent. If I'd really made KCN, that'd have bad implications for us as the ferricyanide entered our slightly alkaline intestines, surely?

But then, when I tested the solution with universal inicator, it was pH7. So that would support my theory that the basic Fe(OH)3 had precipitated out, leaving a neutral solution of KCN. I added more KOH for curiosity's sake. More of the same: more precipitate. This time when I tested the pH, it was pH11-14 (they all look the same to me). That would seem to indicate that all of the ferricyanide had reacted, so the KOH was then just dissolving, yet the solution was still the same yellow colour it started. So does anyone have any idea what happened?
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pHzero
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[*] posted on 29-5-2009 at 13:13


D'oh, I've got it.

The KOH reduces the Fe 3+ to Fe 2+, turning the ferricyanide (K3Fe(CN)6) to the ferrocyanide (K4Fe(CN)6), which is less soluble in water, so some came out of the solution. That would explain why the "preciptitate" was a different colour to Fe3+, and the solution remained yellow.
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[*] posted on 30-5-2009 at 05:44


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%.




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pHzero
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[*] posted on 30-5-2009 at 05:52


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.


Ah I see. It was probably a bad idea to pour it down the sink then, but ah well, it was only a little bit. I should really invest in some H2O2 to oxidise cyanides before I tip them away.
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[*] posted on 30-5-2009 at 07:54


I'm aware of an investigation
http://dx.doi.org/10.1002/zaac.19130840116
With several different outcomes.
Though I have to wonder if your precipitate was an insoluble salt formed from impurities in whatever.
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entropy51
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[*] posted on 30-5-2009 at 11:47


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.
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pHzero
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[*] posted on 30-5-2009 at 11:59


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.


Ooh good call. I'll try that tonight, when my parents have gone to bed. (They get annoyed when I have "chemicals" in the house). I've got some iron (iii) chloride, so if i add some of that and it turns blue, that'll be a positive for ferrocyanide. Yesterday I mixed iron (iii) chloride and ferricyanide and it went mucky brown-green, but then turned a beautiful shade of blue in sunlight as the light reduced the ferricyanide to ferrocyanide, making prussian blue. I even tried some cyanotype printing and managed to form some vaguely recognisable images.
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pHzero
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[*] posted on 30-5-2009 at 16:57


I just did it again and added Fe3+ this time, here's the result:

http://www.youtube.com/watch?v=o--kSFwGTtw

You can see prussian blue forming at the top when i added the Fe3+

http://img200.imageshack.us/img200/4066/1000294.jpg you can see it better in that pic
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[*] posted on 30-5-2009 at 17:14


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.
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[*] posted on 31-5-2009 at 03:09


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.
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[*] posted on 31-5-2009 at 04:59


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]
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pHzero
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[*] posted on 31-5-2009 at 05:13


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]


Hmm, that's a very good point - I really ought to invest in some stoppers and start using the safety goggles my mym bought home from work.
The reason my parents get annoyed though is that I spilt Fe3+ on my carpet a while ago which wont come off (I should try oxalic acid at some point) and made the house smell of acetic acid on two seperate occasions.
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[*] posted on 31-5-2009 at 06:04


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.
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[*] posted on 31-5-2009 at 10:04


Doesn't KCN decompose to HCN in moisture?

It really depends on the concentration of hydroxide. A small concentration would just make your skin a little soapy, a large concentration would give you chemical burns. Either way, you should wear goggles. In this experiment, I'd keep H2O2 around at any matter.

I think that this is in fact a reduction. Make solutions of what you have, then try out the test with equal amounts.
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[*] posted on 31-5-2009 at 11:12


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.


Ah well, you live and learn. Now I know to use goggles with KOH, and that the neutralisation of acetic acid's exothermic enough to boil it.

And Rich_Insane: yes it does slowly hydrolyse, but the KOH in the solution pushes the equilibrium towards the KCN+H2O end and away from the KOH+HCN end. That's why cyanides are a lot safer to handle if you add a little bit of alkaline stabiliser.
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[*] posted on 31-5-2009 at 11:38


pHzero, you said that cyanides are a lot safer to handle if you add a little bit of alkaline stabiliser.

I have doubts about this. Metallic cyanides such as KCN are salts of HCN, which is a weak acid. Salts of weak acids are strong bases.

According to the Merck Index, the pH of a 0.1 N solution of KCN is 11. It is already quite alkaline without adding any base as stabilizer.

I think the point here is to avoid any intentional exposure to any chemical, and to use precautions such as gloves and safety glasses to avoid unintentional contact when the inevitable happens, as it always does sooner or later.
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[*] posted on 1-6-2009 at 13:43


Hmm, I always wonder why people fill their test tubes so much, like you did. I always work with less than 20-25% of the test tube filled, this way you can shake.

What you did is very irresponsible. Shaking conc. KOH without safety glasses... You will be so sorry if you get it in the eye.

Oh, and by the way, IIRC CN(-) solutions are quite easily absorbed through the skin.

Please be safe, if you get an accident, it will be bad for the amateur scientist in general...
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[*] posted on 1-6-2009 at 14:11


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!
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[*] posted on 1-6-2009 at 16:17


Yes some labware is expensive but stoppers are not. You can buy huge bags of stoppers on ebay. If you cannot afford the basics then maybe you need to wait before experimenting.
In the US anyways I have even been able to find rubber and cork stoppers at certain hardware stores. There is no excuse for not having stoppers. Not sure why they have stoppers except to assist those with wine cellars in their basements.

Also, pharmacies carry boxes of disposable latex and nitrile gloves. They rip easy and are not full proof but they are great if you grab a beaker that has potassium permanganate on the outside that you are not aware of. The glove gets it but your hand does not.




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[*] posted on 2-6-2009 at 01:09


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.




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[*] posted on 2-6-2009 at 02:17


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.

Not sure on the details of this reaction, but I would have thought that the only iron to give off an electron would be Fe(II) - if Fe(III) gave off an electron you'd get Fe(IV), not impossible but I'm guessing you didn't mean this?




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[*] posted on 2-6-2009 at 03:29


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

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





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[*] posted on 2-6-2009 at 04:12


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]
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[*] posted on 2-6-2009 at 05:02


Quote: Originally posted by woelen  
Sorry for the mistake. That was one of my unclear moments :P

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

I'm guessing the high-spin (d5) Fe(III) converts to low-spin (d5) Fe(III) by the action of light, which then can easily accommodate an additional electron giving (d6) Fe(II) - which is of course diamagnetic. Does this sound reasonable?




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[*] posted on 2-6-2009 at 16:18


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 :D. I've learned a lot from here.
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