A mixture of solid $Sr{\mathrm{SO}}_4$ and solid ${\mathrm{BaSO}}_4$ is shaken up with water until equilibrium is established. Given that ${\mathrm{K}}_{\mathrm{sp}}\left({\mathrm{SrSO}}_4\right)=7.5\times {10}^{-7},{\mathrm{K}}_{\mathrm{sp}}\left({\mathrm{BaSO}}_4\right)=6\times {10}^{-8}$, calculate the concentrations of ${\mathrm{Sr}}^{2+},{\mathrm{Ba}}^{2+},{\mathrm{SO}}_4^{2-}$ in the solution at equilibrium.

Although it is a violent poison if swallowed, mercury ($\mathrm{II}$) cyanide, $\mathrm{Hg}(\mathrm{CN}{)}_2$ has been used as a typical (skin) antiseptic. $\left({\mathrm{K}}_{\mathrm{sp}}=1.35\times {10}^{-23},\mathrm{Hg}=200\right)$

What is the molar solubility of this salt in pure water?

Although it is a violent poison if swallowed, mercury $\left(\mathrm{II}\right)$ cyanide, $\mathrm{Hg}(\mathrm{CN}{)}_2$ has been used as a a typical (skin) antiseptic. $\left({\mathrm{K}}_{\mathrm{sp}}=1.35\times {10}^{-23},\mathrm{Hg}=200\right)$

How many milligrams of $\mathrm{Hg}(\mathrm{CN}{)}_2$ dissolve per $\mathrm{litre}$ of pure water?

Although it is a violent poison if swallowed, mercury $\left(\mathrm{II}\right)$ cyanide, $\mathrm{Hg}(\mathrm{CN}{)}_2$ has been used as a typical (skin) antiseptic. $\left({\mathrm{K}}_{\mathrm{sp}}=1.35\times {10}^{-23},\mathrm{Hg}=200\right)$

How many milliliters of water are required to dissolve $1.134 \mathrm{g}$ of the salt?

Calculate the solubility of ${\mathrm{CaC}}_2{\mathrm{O}}_4$ in a solution with a fixed ${\mathrm{H}}_3{\mathrm{O}}^{+}$ concentration of ${10}^{-4}\mathrm{M}$. The oxalate containing species $\left({\mathrm{C}}_2{\mathrm{O}}_4^{2-},{\mathrm{HC}}_2{\mathrm{O}}_4^{-},{\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4\right)$ exist in solution and the relevant equilibria are,

${\mathrm{CaC}}_2{\mathrm{O}}_4\left(\mathrm{s}\right)\rightleftharpoons {\mathrm{Ca}}^{2+}+{\mathrm{C}}_2{\mathrm{O}}_4^{2-} {\mathrm{K}}_{\mathrm{sp}}=2.7\times {10}^{-9}$

${\mathrm{H}}_2{\mathrm{C}}_2{\mathrm{O}}_4+{\mathrm{H}}_2\mathrm{O}\rightleftharpoons {\mathrm{H}}_3{\mathrm{O}}^{+}+{\mathrm{HC}}_2{\mathrm{O}}_4^{-} {\mathrm{K}}_1=5\times {10}^{-2}$

${\mathrm{HC}}_2{\mathrm{O}}_4^{-}+{\mathrm{H}}_2\mathrm{O}\rightleftharpoons {\mathrm{H}}_3{\mathrm{O}}^{+}+{\mathrm{C}}_2{\mathrm{O}}_4^{2-} {\mathrm{K}}_2=5\times {10}^{-5}$

$100 \mathrm{mL}$ of $0.1\mathrm{M} {\mathrm{MgCl}}_2$ solution is mixed with $100 \mathrm{mL}$ of $0.2\mathrm{M} \mathrm{NaOH}$ solution. Given for $\mathrm{Mg}(\mathrm{OH}{)}_2 {\mathrm{K}}_{\mathrm{sp}}=12\times {10}^{-12}$. What is the $\mathrm{pH}$ of the resulting solution in nearest possible integers?

$0.2$ millimoles of ${\mathrm{Zn}}^{2+}$ ion is mixed with ${\left({\mathrm{NH}}_4\right)}_2\mathrm{S}$ of molarity $0.02 \mathrm{M}$. The amount of ${\mathrm{Zn}}^{2+}$remains unprecipitated in $20 \mathrm{mL}$ of this solution would be

(Given : ${\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{ZnS}=4\times {10}^{-24}$)