Electrochemistry: describe question

Standard electrode potentials for three half-reactions are given.

$\ce{Fe^{3+}(aq) + e^- <=> Fe^{2+}(aq)}$ $E^\circ = +0.77$ V

$\ce{I2(s) + 2e^- <=> 2I^-(aq)}$ $E^\circ = +0.54$ V

$\ce{Ce^{4+}(aq) + e^- <=> Ce^{3+}(aq)}$ $E^\circ = +1.44$ V

(a) Define the terms standard electrode potential and standard cell potential.

(b) Describe the standard hydrogen electrode and how it can be used to measure the standard electrode potential of the $\ce{Fe^{3+}/Fe^{2+}}$ half-cell.

(c) When a cell is constructed by connecting the $\ce{Fe^{3+}/Fe^{2+}}$ half-cell and the $\ce{I2/I^-}$ half-cell, deduce the polarity of each electrode and state the direction of electron flow in the external circuit.

(d) Construct the overall redox equation for the reaction that occurs in the cell in (c).

(e) Predict whether $\ce{Ce^{4+}(aq)}$ ions can oxidise $\ce{I^- (aq)}$ ions under standard conditions. Use the data to justify your answer and state which is the stronger oxidising agent.