Potassium Ferricyanide from ferrocyanide via H₂O₂, NaNO₂

I have a lot of potassium ferrocyanide, yellow crystals, but they are generally too stable to electroplate with in non-aqueous solutions. It's hard to get it to ionize, and also difficult to get it to separate from water of hydration.

I've been told the ferricyanide is much less stable and doesn't need water of hydration when crystallized.

I assume I can oxidize ferrocyanide to ferricyanide.
Chlorine gas will do it.

But, I didn't have any on hand today, so I tried something different.
If this new idea isn't a good way to go, I can get calcium hypoclhorite and mix it with sodium bisulfite and table salt to produce chlorine gas; but I'd rather avoid chlorine gas if I can.

Therefore:
I mixed posassium ferrocyanide with water to make a yellow solution. Then I added very weak (1/2 %) hydrogen peroxide. Because --- Nothing appeared to happen. I then I added sodium nitrite ( about the same amount of crystals, 1 [cc] or so as the ferrocyanide).

Very slowly bubbles began forming, and the solution changed from a pure yellow to a yellow-green color over a period of an hour.

I suspect the reaction worked, although the color isn't red yet.
I hope If I allow this greenish solution to evaporate over a period of days, the color will redden and eventually begin to crystallize.

Am I anywhere close to correct ??

I'm not sure if I should expect the sodium nitrite, or peroxy nitrite, or possibly nitrate (if nitrite oxidized), to co-crystalize with the ferricyanide.

How can I optimize the process to get as much ferricyanide as possible with minimum water crystallization ?

N.B.
There is some mention in the literature that a side reaction might happen, which would replace some cyanide with oxygen-nitrate compounds or possibly cyaniamide. If this makes the iron cyanide more ionizable, I don't mind. But, I'd rather not have free oxygen able to combine with iron to make rust.

Thoughts?

Quote: Originally posted by semiconductive

But, I didn't have any on hand today, so I tried something different. If this new idea isn't a good way to go, I can get calcium hypoclhorite and mix it with sodium bisulfite and table salt to produce chlorine gas; but I'd rather avoid chlorine gas if I can.

@Semiconductive, it might be cheaper just to buy it ($23-$36/lb.), unless of course if the goal is the fun of the synthesis, which is great too!

Anyway, here are a couple of US-based sellers. I haven't vetted all of these, so it's at your own discretion.
https://us.chemicalstore.com/potassium-ferricyani995potassiu...
https://stores.photoformulary.com/potassium-ferricyanide/
https://artcraftchemicals.com/products/potassium-ferricyanid...

I'm not sure what I'm doing wrong.

Next attempt, incrementally attempting to change using ammonium chloride.

I mixed yellow prussate crystals with a small amount <10% by volume Ammonium chloride. Then added hydrogen peroxide to dissove it (~<3%).

The result is much the same as before, I get a brighter yellow-green color.

Therefore, I continued to add Ammonium Chloride and dissolve it until I had added around 25% crystals by volume compared to ferrocyanide. It became brighter yellow, and then the green became much more saturated but otherwise ... still green. It's not turning red, or reddish orange.



So, I dried it overnight with a gentle air breeze in the test-tube. It mostly dries super bright yellow (more yellow than the original salt) with a small part on top being yellowish orange. But, all references I see say the solid ought to be "ruby red" -- and it is not.

Mixing it with weak peroxide, again, I added more ammonium chloride. The green starts to go dark, and the more ammonium chloride I add, the darker the green color is. The solution itself looks almost black, which could be red -- but if it is, then why isn't it drying red rather than turning orange as it evaporates ?



If I had stripped of the potassium, completely, I would be left with prussian blue and that's not what I have. I would need at least dark green for it to be a mix of prussian blue with ferrocyanide.

I estimate close to 50% by volume is ferrocyanide vs. ammonium chloride, now.

Drying the solution under air causes whitish-yellow crystals to separate out first and fall to the bottom of the test tube, with orange crystals forming on the test tube walls above the salt. Orange is not ruby red.

The more ammonium chloride I add, the more orange salt appears -- but the yellow stuff is still there. Heating it makes the crystals falling to the bottom of the test tube become a dark green. I assume that's the onset of prussian blue....

The light orange makes me suspect that I have half yellow, half red, iron cyanides mixed ?

Apparently, it's difficult to drive the reaction to a pure ferric state ?

Online searches don't give me any clues, either.

Anybody have an idea of what I'm doing wrong?
(Besides not buying the ferric already, which I am ordering...)
This is supposed to be an *easy* reaction to do.






[Edited on 27-3-2026 by semiconductive]

problems with posting....

but:

"Oxygen exerts no action on ferrocyanide of potassium, but
ozone
electric current
bromine
diluted nitric acid
peroxide of lead
manganese dioxide
potassium permanganate
all transform it either wholly or in part into potassium ferricyanide."

page 29. "The Cyanide industry - theoriticallly and practically considered. by R Robine M Lenglen 1906

PLEASE BE CAREFULL - on using acids on ferrocyanide - READ before proceding.

here is a table from bellow:

https://www.911metallurgist.com/blog/ferrocyanides/

[Edited on 27-3-2026 by RU_KLO]

[Edited on 27-3-2026 by RU_KLO]

Indeed. I have avoided using any strong acids.
I have No death wish for people around me.


The 911 metals site notes:

How strange!
Bright yellow, then, is sometimes ferricyanide and not mixed ferro-ferri cyanides?!


I may have already succeeded if H₂O₂ has the same effect as acetic acid with sodium nitrite.

There's no guideline as to how concentrated the acetic acid is, either, ... but I might try replicating it by slow addition (dropwise) of a weak solution first, then glacial later (outside with ice).

911met.. also mention no precipitation when mixed with caffiene. But, they don't say if that's because there's no reaction -- or if it forms a salt. Caffiene even dies a sluggish death in cyanide.

A lot of online references say that Chlorine conversion is not very high yield, but the halogens mentioned by 911metallurigist say "quantitative" even if slow. eg: A small amount of Iodine will produce ferricyanide + a small amount of prussian blue precipitate. ( But: I'm temporarily out of iodine crystals. Used them to very effectively kill off a toe nail fungus with tea tree oil / eucalyptis .... )

There seems to be quite a bit of contradiction about what colors are supposed to appear...


[Edited on 27-3-2026 by semiconductive]

That (A.CN) should be a clear solid mixed with light green iron chloride crystals, and this reaction is close to the experiment that I tried.

It confirms my thinking that chloride can displace potassium in yellow prussate. But: Chloride can also displace the cyanide... at least when nearly dry?! hmm... I wonder why that is.

HCl is a strong acid, and therefore it must have a weak bond that is easily broken.

The weaker the acid, generally, the stronger the bond. Cyanides, therefore, should have strong bonds with small amounts of ionization because HCN is a very weak acid, even weaker than H₂CO₃.

But under heat there is a non-intuitive situation where a likely strong bond (Fe-CNx) is replaced by a weaker bond -- or maybe, the cyanide likes ammonia even more than it likes iron and leaves iron ?

P. 92 ...
Ferrocyanides come in two colors, lemon yellow and amber. According to S.H.C. Briggs, it's because there are two "stereoisomerides" that can be formed. The amber is formed when 1% acetic (or other acid) is added and kept out of contact with air for 24 hours. Adding alcohol then precipitates a cream colored material.

That is verrrry interesting.
It gives me a plausible explanation why the orange-ish crystals dry mostly yellow as air gets them. But why does alcohol make it turn 'cream'?

But I think I can test it!!!!

0.1 [cc] of crystals, 4 [cc] H2O dissolves them, 2 [cc] of paraffin/kerosene to cap it off.
Use soldering iron which extremely slowly will boil water out and cause and dissolved oxygen/nitrogen to go away. The solution is almost clear, just the tiniest tinge of dull yellow.

I take a clean glass stir rod .. touch to the top of glacial acetic acid with it. Shake off droplet so only a tiny film remains on glass ... and push it through into the bubbly hot water.

Instant yellow, quite visible, and it's turning rapidly yellow green. But it's not bright yellow like with hydrogen peroxide or NaNitrite. I think it likes to make a hot coloidal solution even with tiniest amounts of acetic acidifying.



Now, let's see how it crystallizes (eventually).



The colloidal particles fall out just above the soldering iron, which is maintaining the glass bottom at 150[°C]. Pretty much there isn't enough surface area to transfer heat effectively, so we get microboiling right at the point of contact with the glass. When I first heat the test tube, gas bubbles are released and the tube gets hot to the touch (but not enough to burn). But after a half hour or so, no more gas bubbles appear and all semblance of boiling stops.

Basically, a cycle is set up where coloidal particles (white to the eye, or a paler yellow than the liquid -- though the picture has them in shadow) build up above the hottest part of the test tube. Then they slide down onto the hot spot, and explode into dispersed particles again with a tiny bubble or two of gas escaping. Then the glass is clean ... and over a period of an hour the colloidal particles build up again.

There's two ways I expect this experiment might progress. One, the acetate may slowly be gasiffied and removed from solution because the cyanide is more stable (AKA: a weaker acid). In that case the solution ought to become more and more yellow, and closer to clear water. Or the acetate might stay at a constant value because it's made a salt that's stable in this environment.

Depending on how it progresses in the next few days, I might need to add more acetate solution; or I might need to think of a different way to introduce acetate in smaller amounts.

At first it looked like it was becoming darker green with time, but now it looks to be turning a lighter brighter yellow (slowly).

This is going to take a few days ...
The ferrocyanide is reagant grade tri-hydrate from Biopharm™.


-------------------------------

P. 97 -- calcium ferrocyanide, formed from CaCl₂ + NH₄Cl + sodium ferrocyanide.
Hm... maybe I could make that with potassium ferrocyanide. I wonder.

Next p. 97..98 says "Calcium ferrocyanide is a very soluble salt. ... the clear concentrated liquor remaining uncrystallized for a long time."

But, Other sources, say it isn't soluble in water (2026):
https://en.wikipedia.org/wiki/Calcium_ferrocyanide

So, that confirms my suspicion about color -- intense yellow is what a ferricyanide solution is normally expected to be when not saturated. Saturated solutions are red.

The fact that it's a powerful reducing agent in the presence of caustic alkalies gives me pause. Reducing agents are usually oxidized by atmospheric oxygen. But, if ferricyanide is converted to ferrocyanide by H₂O₂, that's a reduction of the iron state. Very ... confusing.

Can oxygen in air can reduce ferricyanide to ferrocyanide in alkali conditions as well?

-------------------------

Over-all, "The Chemistry of Cyanogens" was a very satisfying read. Quite a bit of useful information in it. Thank you, @Boffis !

I got the impression that thiocyanates can be made from inexpensively from ferrocyanide + sulfur (elemental). Thicyanates are soluble in both water and alcohol, and don't have a stubborn water of crystalization.

The same might be done with a selenide or telluride, but those are much less stable for some reason.

On the opposite side of the spectrum, neutralization/remediation of cyanide poisoning is done with non-toxic hypo. Sodium Thiosulfate. Therefore I will be getting some as an extra safety precaution with ferricyanide.

[Edited on 28-3-2026 by semiconductive]

Quote: Originally posted by DraconicAcid

Quote: Originally posted by semiconductive   But, I didn't have any on hand today, so I tried something different. If this new idea isn't a good way to go, I can get calcium hypoclhorite and mix it with sodium bisulfite and table salt to produce chlorine gas; but I'd rather avoid chlorine gas if I can. I don't have any experience with ferrocyanides or ferrocyanides, but hypochlorites with react with bisulphAtes to give chlorine gas. With sulphItes or bisulphItes, you get sulphates and chloride ions. I'm assuming that's just a typo on your part, but just making sure.

You're not wrong. I have avoided chlorine gas, correct?

I also think I may have found another substance that might be capable of electrodepositing iron pyrite. eg:The antidote to cyanide poisioning. (Edit ThiosulfAte is an awesome suspect especially the form it becomes after it has reduced ferric ions.

Without a polysulfide present, it might or not work; but I think there is an excellent chance it will still give me either FeS or FeS₂.

Do you disagree?


[Edited on 28-3-2026 by semiconductive]

Pictures of results of mild acidification of ferrocyanide with only a dipped and shaken off stir rod's film worth of acetic acid.

eg: After 8 hours after last picture, but just before mild vacuum was pulled. The orange deposit on the glass does not slide down to bottom of tube as did settling colloid from 8 hours ago.



And 5 minutes after pulling a mild (aspirator) vaccuum via a flitr8 pump.



Over the next 48 hours, the orange colloidal particles will settle out again. The solution will go back to looking like the brighter yellow-green shade that is shown in the first of the two pictures. This is only slightly more intense than the originally dissolved ferrocyanide's pale yellow color.