Define what is meant by standard cell potential.
Draw a fully labelled diagram of the experimental arrangement that could be used to measure the standard electrode potential of the $\text{Pb}^{2+}\text{(aq)}/\text{Pb(s)}$ electrode. Include the chemicals needed.
The $E^{\circ}$ for a $\text{Pb}^{2+}\text{(aq)}/\text{Pb(s)}$ electrode is $-0.13\,\text{V}$. Suggest how the $E$ for this electrode would change if the concentration of $\text{Pb}^{2+}\text{(aq)}$ ions were lowered. Show this by putting a tick ($\checkmark$) in the right box. Explain your response.
Car batteries contain rechargeable lead-acid cells. In each cell, the negative electrode is Pb metal and the positive electrode is $\text{PbO}_2$. The electrolyte is $\text{H}_2\text{SO}_4\text{(aq)}$. During discharge of a lead-acid cell, $\text{Pb}^{2+}$ ions are deposited as $\text{PbSO}_4\text{(s)}$ at the negative electrode. $\text{Pb(s)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{PbSO}_4\text{(s)} + 2\text{e}^- $ Calculate the mass of Pb that turns into $\text{PbSO}_4$ when the cell supplies a current of $0.40\,\text{A}$ for $80\,\text{minutes}$.
Complete the half-equation for the reaction that takes place at the positive electrode. $\text{PbO}_2\text{(s)} + \text{SO}_4^{2-}\text{(aq)} + \ldots + \ldots \rightarrow \text{PbSO}_4\text{(s)} + \ldots$
The diagrams show how the voltage across two different cells changes with time when each cell is used to provide an electric current. One diagram is for a lead-acid cell and the other is for a $\text{H}_2/\text{O}_2$ fuel cell. Suggest a reason why the voltage of the lead-acid cell changes after several hours, and why the voltage of the fuel cell remains constant.