1
$\begingroup$

This patent notes in passing:

It may be necessary to add $\ce{NaOH}$ to stabilize aqueous $\ce{NaCN}$. Anyone skilled in the art would know the amount of $\ce{NaOH}$ to add.

Obviously I'm not skilled in the art. I assume the stabilization is against hydrolysis. What is the mechanism by which $\ce{NaOH}$ would stabilize aqueous $\ce{NaCN}$, and how would one determine the correct amount of $\ce{NaOH}$ to add?

(I'm expecting it's some target pH level, but I'm too far along my "assumption" chain to know what the issues are.)

|improve this question|||||
$\endgroup$
  • 2
    $\begingroup$ The mechanism is that the $\ce{OH^{-}}$ would drive those $\ce{HCN}$ (resulting from hydrolysis) all the way back to $\ce{CN^{-}}$. As for the amount... well, yes, looks mighty like it has to do with pH. $\endgroup$ – Ivan Neretin Jan 15 '16 at 17:25
  • $\begingroup$ The pH is the tip of the ice berg. Many other safety protocols should be observed if working with cyanide solutions. What protocols would depend on the reaction, the nature of the setup (e.g. glassware or apparatus), and the amount of cyanide involved. $\endgroup$ – MaxW Jan 15 '16 at 19:59
5
$\begingroup$

Hydrogen cyanide is an acid with a pKa of 9.2, meaning that at pH 9.2 in water you have about a 50-50 mix of cyanide anion and hydrogen cyanide. If you used sodium hydroxide to adjust the pH of a NaCN solution to, say, 12.2, you would inhibit the decomposition of NaCN by shifting the equilibrium to 99.9 percent cyanide anion and 0.1% hydrogen cyanide. Each pH unit higher gives you an additional factor of 10 lower concentration of HCN.

Another factor is temperature: since HCN boils at 28 C, it tends to evaporate from a room-temperature solution of NaCN, driving the reaction to the right. The solution would be more stable if kept cold.

Finally, it would be better to keep the solution in a tightly closed bottle, since carbon dioxide could be absorbed from the air, acidifying the solution and releasing hydrogen cyanide.

|improve this answer|||||
$\endgroup$
  • $\begingroup$ So (in a closed container) you can almost eliminate hydrolysis of a moderate acid using a strong base? $\endgroup$ – feetwet Jan 15 '16 at 22:27

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.