i) State the chemical composition of solution A, electrode B. ii) Complete the diagram so that the full experimental set-up is shown.
The cell shown above is not under standard conditions because the $[\text{Ag}^+]$ in a saturated $\text{AgCl}$ solution is much smaller than $1.0\ \text{mol dm}^{-3}$. The $E_{electrode}$ is linked to $[\text{Ag}^+]$ by the equation below. $E_{electrode} = E_{electrode}^\circ + 0.06\log [\text{Ag}^+]$ i) Use the Data Booklet to work out the $E_{cell}^\circ$ if the cell were operating under standard conditions. In the experiment described above, $E_{cell}$ was measured as $+0.17\ \text{V}$. ii) Calculate $E_{electrode}$ for the $\text{Ag}^+/\text{Ag}$ electrode in this experiment. iii) Use equation 1 to determine $[\text{Ag}^+]$ in the saturated solution.
State the composition of solution A, electrode B.
Complete the diagram so that the full experimental set-up is shown.
Use the Data Booklet to work out the $E^\circ_{\text{cell}}$ if the cell were operating under standard conditions.
Calculate $E_{\text{electrode}}$ for the $\text{Ag}^+/\text{Ag}$ electrode in this experiment.
Use equation 1 to determine $[\text{Ag}^+]$ in the saturated solution.
Write an expression for the $K_{sp}$ of silver sulfate, $\text{Ag}_2\text{SO}_4$, and include units.
Calculate the value of $K_{sp}$ for silver sulfate.
Describe how the colours of the silver halides, together with their relative solubilities in $\text{NH}_3(\text{aq})$, can be used to tell apart solutions containing the halide ions $\text{Cl}^-$, $\text{Br}^-$ and $\text{I}^-$.
Describe and explain the trend in the solubilities of the sulfates of the Group II elements.