Thus far we have been picking exclusively on reactions involving CN\". but there

Question

Thus far we have been picking exclusively on reactions involving CN". but there s no reason why we shouldn't give equal opportunity to Ag^+. For example, you're probably aware that some metal ions, like Ag^+, form insoluble hydroxides. Furthermore this process is going to be tough to avoid in strongly basic solutions. So consider the reaction Ag_2O(s) + H_2O 2Ag^+ 2OH^- K_sp = 3.8 times 10^-16 If Ag_2O forms to an appreciable extent what effect will this equilibrium have on the solubility of AgCN? Explain your reasoning. Judging by the equilibrium reaction shown, would you expect the effect of this equilibrium on the solubility of AgCN to be pH dependent? Try this little experiment. Pick the solution from Table 1 where your reasoning suggests that the precipitation of Ag_2O will be most significant. You can easily figure out [OH^-] from the pH and [Ag^+] that is given you, so calculate the ion product suggested by the reaction above and determine whether or not precipitation of Ag_2O will actually occur.

Explanation / Answer

In this case the Ag ion would be a common ion and therefore it is a competitive ion the both reactions, by L'chatelier principle if the Ag2O reaction shifts to the left and more Ag2O is formed this reaction requires more Ag ion concentration to maintain the Kps, then the effect on the AgCN equilibrium would be to shift to the right to maintain that concentration of Ag ions product of both equilibrium, therefore the solubility of AgCN would increase, base on the Ag2O reaction it would be pH dependent because pH is an indirect measure of OH related by pOH for the equation: pH + pOH = 14.

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