How will the concentration of ${\mathrm{Ag}}^{+}$ in a saturated solution of $\mathrm{AgCl}$ diminish if such an amount of $\mathrm{HCl}$ is added to it that the concentration of the ${\mathrm{Cl}}^{-}$ in the solution becomes $0.03 \mathrm{mol}/\mathrm{L}$? ${\mathrm{K}}_{\mathrm{sp}}\left(\mathrm{AgCl}\right)=1.8\times {10}^{-10}$.

Given that $2\times {10}^{-4} \mathrm{mol}$ each of ${\mathrm{Mn}}^{2+}$ and ${\mathrm{Cu}}^{2+}$ were contained in one litre of a $0.003 \mathrm{M} {\mathrm{HClO}}_4$ solution and this solution was saturated with ${\mathrm{H}}_2\mathrm{S}$. Determine whether each of these ions, ${\mathrm{Mn}}^{2+}$ and ${\mathrm{Cu}}^{2+}$, will precipitate as a sulphide or not. The solubility of ${\mathrm{H}}_2\mathrm{S}$,  $0.1 \mathrm{mol}/\mathrm{L}$, is assumed to be independent of the presence of other materials in the solution.

${\mathrm{K}}_{\mathrm{sp}} \left(\mathrm{MnS}\right)=3\times {10}^{-14}, {\mathrm{K}}_{\mathrm{sp}} \left(\mathrm{CuS}\right)=8\times {10}^{-37}$

${\mathrm{K}}_1$ and ${\mathrm{K}}_2$ for ${\mathrm{H}}_2\mathrm{S}$ are $1\times {10}^{-7}$ and $1.1\times {10}^{-14}$, respectively. Also, calculate the percentage of $\mathrm{Cu}$ left unprecipitated. Will $\mathrm{MnS}$ precipitate if the above solution is made neutral by lowering $\left[{\mathrm{H}}^{+}\right]$ to ${10}^{-7} \mathrm{M}$?

What $\mathrm{pH}$ must be maintained in a solution saturated in ${\mathrm{H}}_2\mathrm{S} (0.1 \mathrm{M})$ and ${\mathrm{Zn}}^{2+}$ $\left({10}^{-3} \mathrm{M}\right)$ to prevent $\mathrm{ZnS}$ from precipitating?

${\mathrm{K}}_{\mathrm{sp}}\left(\mathrm{ZnS}\right)=1\times {10}^{-21}, {\mathrm{K}}_{\mathrm{a} }\left({\mathrm{H}}_2\mathrm{S}\right)=1.1\times {10}^{-21}$

Will $\mathrm{FeS}$ precipitate from a solution that is saturated in ${\mathrm{H}}_2\mathrm{S} (0.1 \mathrm{M})$ and ${\mathrm{Fe}}^{2+} \left(0.002 \mathrm{M}\right)$ at $\mathrm{pH}=3.5$?

${\mathrm{K}}_{\mathrm{sp}} \left(\mathrm{FeS}\right)=6.3\times {10}^{-18}, {\mathrm{K}}_{\mathrm{a} }\left({\mathrm{H}}_2\mathrm{S}\right)=1.1\times {10}^{-21}$.

A buffer solution is $0.25 \mathrm{M} {\mathrm{CH}}_3\mathrm{COOH}, 0.15 \mathrm{M} {\mathrm{CH}}_3\mathrm{COONa}$, saturated in ${\mathrm{H}}_2\mathrm{S}$ $(0.1 \mathrm{M})$ and has $\left[{\mathrm{Mn}}^{2+}\right]=0.015 \mathrm{M}$. 

${\mathrm{K}}_{\mathrm{a}} \left({\mathrm{CH}}_3\mathrm{COOH}\right)=1.74\times {10}^{-5}$, ${\mathrm{K}}_{\mathrm{a}} \left({\mathrm{H}}_2\mathrm{S}\right)=1.1\times {10}^{-21}$ and ${\mathrm{K}}_{\mathrm{sp}} \left(\mathrm{MnS}\right)=2.5\times {10}^{-13}$.

Will $\mathrm{MnS}$ precipitate?

A buffer solution is $0.25 \mathrm{M} {\mathrm{CH}}_3\mathrm{COOH}-0.15 \mathrm{M} {\mathrm{CH}}_3\mathrm{COONa}$, saturated in ${\mathrm{H}}_2\mathrm{S}$ $(0.1 \mathrm{M})$ and has $\left[{\mathrm{Mn}}^{2+}\right]=0.015 \mathrm{M}$. 

${\mathrm{K}}_{\mathrm{a}}\left({\mathrm{CH}}_3\mathrm{COOH}\right)=1.74\times {10}^{-5}$, ${\mathrm{K}}_{\mathrm{a}}\left({\mathrm{H}}_2\mathrm{S}\right)=1.1\times {10}^{-21}$ and ${\mathrm{K}}_{\mathrm{sp}}\left(\mathrm{MnS}\right)=2.5\times {10}^{-13}$

Which buffer component should be increased in concentration and to what minimum value to just start the precipitation of $\mathrm{MnS}$?