It has been a long time since I studied organic chemistry, but one thing I do remember is that when we needed cyanide salts we were told not to bother keeping any remaining solution because the salts hydrolyze very quickly.

But can't hydrolyzed salts be fully recovered by simply evaporating the water?

E.g., we don't worry about evaporating NaCl from a distilled water solution and finding anything other than NaCl. So I'm wondering if the lab instructions to neutralize and dispose of cyanide salt solution was just based on expediency. Or: can something unsafe happen during the evaporation of a hydrolyzed salt solution?

  • $\begingroup$ Do you mean "hydrolyzed", or just "aqueous" or "absorbed water"? When I hear "hydrolyzed" in a chemistry context, I think of covalent bonds being broken, but I can't work out if that is what you mean... $\endgroup$
    – Curt F.
    Commented Jul 23, 2015 at 20:31
  • $\begingroup$ @CurtF. -- Again, I'm rusty, but I'm pretty sure I mean hydrolyzed. If it was just aqueous then you just dry it out and everything's fine. But my recollection of the reason we don't store cyanide salts in solution is that they actually decompose via hydrolysis, e.g., KCN(s) -> K(+aq) + CN(-aq). Which affects the solution's strength as a reagent, I guess? $\endgroup$
    – feetwet
    Commented Jul 23, 2015 at 20:44
  • $\begingroup$ That is just dissolution. In fact if you have KCN "in solution", that means it consists of K(+)(aq) and CN(-)(aq) ions dissolved in the solution. Maybe you are talking about hydrolysis of cyanide complexes of transition metals, such as ferrocyanide or ferricyanide? $\endgroup$
    – Curt F.
    Commented Jul 23, 2015 at 21:06
  • $\begingroup$ @austinian: Oh, that might explain it. So cyanide salts are hydrolyzed into hydrogen cyanide? That's a good reason not to leave the solution sitting around! $\endgroup$
    – feetwet
    Commented Jul 23, 2015 at 21:06
  • $\begingroup$ No they are not hydrolyzed to hydrogen cyanide. OH- is a very strong base and HCN is an acid, so if any such reaction formed they would immediately back convert to CN- and H2O. Perhaps if you had cyanide salts in open containers, over long periods of time (several days or weeks) enough HCN would evaporate to significantly lower CN- concentrations in the solution. But that won't happen if you store in sealed containers. $\endgroup$
    – Curt F.
    Commented Jul 23, 2015 at 21:08

1 Answer 1


Simple cyanide salts

In aqueous solution, simple cyanide salts more or less completely dissociate into their constituent ions. E.g. for potassium cyanide:

$\ce{{KCN}(s) + H2O(l) -> K+(aq) + CN- (aq)}$

Those solutions are relatively stable. There is an acid-base equilibrium between $\ce{CN-}$ and $\ce{HCN}$, but in unbuffered or in alkaline aqueous solution this equilibrium will be almost completely on the left:

$\ce{CN- + H2O <<=> HCN + OH-}$

Over very long times, if the solution is left in an open container, some of the tiny, minute amount of $\ce{HCN}$ formed in this way could evaporate, leading to a very slow loss of cyanide from the solution. If stored closed this will not be a problem.

(Mixing simple cyanide salts with an acid will result in an immediate and dangerous release of gaseous $\ce{HCN}$, so in this question I assume that no acids are involved.)

Thus, for simple cyanide salt solutions stored appropriately, yes you should indeed be able to recover all of the initially added the solid salt after boiling off or otherwise removing any water.

Cyanide complexes of transition metals

Salts such as potassium ferricyanide $\ce{K3Fe(CN)6}$ consist of potassium cations and complexed ferricyanide anions, $\ce{Fe(CN)6^{~3-}}$. They are soluble in water as well, but unlike their simple cyanide anions, they can be hydrolyzed.

The hydrolysis is catalyzed by UV light and results in the formation of free $\ce{CN-}$ anions as well as iron hydr(oxide) solid precipitates.

See this book for more information.

  • $\begingroup$ Thanks again for an outstanding answer. I just had to see the photolytic reaction you referenced, and that has led to some more questions. When you have the time could you take a look at those at chemistry.stackexchange.com/q/34398/5354 ? $\endgroup$
    – feetwet
    Commented Jul 26, 2015 at 20:46

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