# Why is silver nitrate used for cyanide titration?

Or more to the point, since silver isn't exactly cheap: Why can't cuprous nitrate ($\ce{CuNO3}$) substitute for the pricier ($\ce{AgNO3}$) in cyanide titration?

I assume that free cyanides bind eagerly with all Group 9-11 metals, since they're used industrially to leach the heavier ones.

I'm guessing that there's something relatively quirky about the Period 4 metals Co, Ni, and Cu that prevents their use as a titrant of aqueous cyanide: E.g., it appears that nickel nitrates are very hygroscopic.

• Out of curiosity (it kind of affects my answer), does this work with $\ce{HCN}$? Or just the $\ce{CN^-}$ ion? Feb 23 '16 at 23:09
• @DGS - AFAIK cyan salts (which I would assumed to include HCN) are always titrated in aqueous solution. However the pH of the solution is not specified, which leaves me wondering whether it is assumed to be neutral or whether that is irrelevant to the titration.... Feb 23 '16 at 23:23
• @DGS - After further thought: The titration is for "free cyanide." My guess is that CN ions that have bound to H in solution are not "free" and won't bind to other metals. (Therefore, if someone is titrating for cyanide reagent quality they probably want to do it in a very alkaline solution?) In fact, hydrolysis of cyan salts is notorious for producing HCN (en route to the usual irreversible nitrile hydrolysis). Feb 24 '16 at 14:27
• In that case, I'm sticking with what I've already said below. And yes, they would have to adjust the pH to significantlt higher than 9.2 (pKa of HCN in water). Cool question. Feb 24 '16 at 16:47

Due to the hardness of copper ions relative to silver ions, copper ions will bind hydroxide more readily in the presence of cyanide than silver ions do. Once ligated by $\ce{OH^-}$ ions, $\ce{Cu^{1+}}$ will oxidize to $\ce{Cu}^{2+}.$
The softer acid, $\ce{Ag^{1+}}$, doesn't do this nearly as quickly, especially not in the presence of a softer ligand like cyanide.