(c) Car batteries are made up of rechargeable lead-acid cells. Each cell consists of a negative electrode made of Pb metal and a positive electrode made of PbO₂. The electrolyte is H₂SO₄(aq). When a lead-acid cell is in use, Pb²⁺ ions are precipitated out as PbSO₄(s) at the negative electrode. Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻ (d) The diagrams show how the voltage across two different cells changes with time when each cell is used to provide an electric current. [Figure: Two graphs showing Voltage/V vs time/hours. One for a lead-acid cell, showing decreasing voltage over time. One for an H₂/O₂ fuel cell, showing constant voltage over time.]

1. (a)Define the term standard cell potential.[2]
2. (b)(i)Draw a fully labelled diagram of the experimental set-up you could use to measure the standard electrode potential of the Pb²⁺(aq)/Pb(s) electrode. Include the necessary chemicals.[4]
3. (b)(ii)The E for a Pb²⁺(aq)/Pb(s) electrode is -0.13V. Suggest how the E for this electrode would differ from its E⁰ value if the concentration of Pb²⁺(aq) ions is reduced. Indicate this by placing a tick (√) in the appropriate box in the table. more negative | no change | less negative ---|---|--- Explain your answer.[2]
4. (c)(i)Calculate the mass of Pb that is converted to PbSO₄ when a current of 0.40A is delivered by the cell for 80 minutes.[2]
5. (c)(ii)Complete the half-equation for the reaction taking place at the positive electrode. PbO₂(s) + SO₄²⁻(aq) + ______ + ______ → PbSO₄(s) + ______[1]
6. (d)Suggest a reason why • the voltage of the lead-acid cell changes after several hours, • the voltage of the fuel cell remains constant.[2]

This question uses a diagram from the original paper, which isn't shown here. Open the full paper from the past papers archive to see it.