Write the solubility product expression, $K_{sp}$, for $\text{Ag}_2\text{SO}_4$ and give its units.
Calculate the value of $K_{sp}(\text{Ag}_2\text{SO}_4)$ at $298\,\text{K}$.
Using $\text{Ag}_2\text{SO}_4$ as the example, complete the Hess' Law energy cycle that links lattice energy, $\Delta H^{\circ}_{latt}$, enthalpy change of solution, $\Delta H^{\circ}_{sol}$, and enthalpy change of hydration, $\Delta H^{\circ}_{hyd}$. On the diagram, place the correct species in the two blank boxes; add the proper symbol to each enthalpy change; and finish the two missing arrows so that the directions of enthalpy change are shown correctly.
Use the Data Booklet to find $E^{\circ}_{cell}$ under standard conditions and state which electrode is positive.
How does the actual $E_{cell}$ of the cell above compare with $E^{\circ}_{cell}$ under standard conditions? Give a reason for your answer.
How would $E_{cell}$ of the cell above change, if at all, if a few $\text{cm}^3$ of concentrated $\text{Na}_2\text{SO}_4(aq)$ were added to the beaker containing $\text{Fe}^{3+}(aq)$ and $\text{Fe}^{2+}(aq)$, and to the beaker containing $\text{Ag}_2\text{SO}_4(aq)$?
Explain any change in $E_{cell}$ that you stated in (iii).
Iron(III) sulfate solutions are acidic because of the following equilibrium: $[\text{Fe(H}_2\text{O)}_6]^{3+}(aq) \rightleftharpoons [\text{Fe(H}_2\text{O)}_5(\text{OH})]^{2+}(aq) + \text{H}^+(aq)$, $K_a = 8.9 \times 10^{-4}\,\text{mol dm}^{-3}$. Determine the $\text{pH}$ of a $0.1\,\text{mol dm}^{-3}$ solution of iron(III) sulfate, $\text{Fe}_2(\text{SO}_4)_3$.