State the chemical composition of solution A and electrode B.
Complete the diagram so that the entire experimental arrangement is shown.
The cell above is not under standard conditions because the [$\text{Ag}^+]$ in a saturated solution of $\text{AgCl}$ is far below $1.0\,\text{mol dm}^{-3}$. The $E_{\text{electrode}}$ is linked to [$\text{Ag}^+]$ by the equation below: $E_{\text{electrode}} = E^\circ_{\text{electrode}} + 0.06\,\log[\text{Ag}^+]$ Use the Data Booklet to work out $E^\circ_{\text{cell}}$ if the cell were operating under standard conditions.
In the experiment above, $E_{\text{cell}}$ was recorded as $+0.17\,\text{V}$. Calculate $E_{\text{electrode}}$ for the $\text{Ag}^+/\text{Ag}$ electrode in this experiment.
Use equation 1 to calculate [$\text{Ag}^+$] in the saturated solution.
Write an expression for $K_{sp}$ of silver sulfate, $\text{Ag}_2\text{SO}_4$, and include units.
Using a similar experimental arrangement to the one shown opposite, it is found that $[\text{Ag}^+]$ in a saturated solution of $\text{Ag}_2\text{SO}_4$ is $1.6 \times 10^{-2}\,\text{mol dm}^{-3}$. Calculate the value of $K_{sp}$ of silver sulfate.
Describe how the colours of the silver halides, together with their relative solubilities in $\text{NH}_3\,(aq)$, may be used to tell apart solutions containing the halide ions $\text{Cl}^-$, $\text{Br}^-$ and $\text{I}^-.$
Describe and explain the pattern in the solubilities of the sulfates of the elements in Group II.