Complete the diagram so that it shows a standard hydrogen electrode. Add labels to your diagram. Identify every substance. You do not need to state the standard conditions.
An electrochemical cell is assembled with an $\text{Fe}^{3+}/\text{Fe}^{2+}$ electrode and a standard hydrogen electrode. Identify the positive electrode in the electrochemical cell and the direction in which electrons travel in the external circuit.
The vanadium-containing species in the electrode reactions listed in Table 3.1 are V, $\text{V}^{2+}$, $\text{V}^{3+}$, $\text{VO}^{2+}$ and $\text{VO}_2^+$. Identify one vanadium-containing species that does not react with $\text{Fe}^{2+}$ ions under standard conditions. Use data from Table 3.1 to explain your answer.
Identify all the vanadium-containing species that react with $\text{Fe}^{2+}$ ions under standard conditions.
Write an equation for one of the possible reactions identified in (ii).
A second electrochemical cell is assembled with an $\text{Fe}^{3+}/\text{Fe}^{2+}$ electrode and an alkaline $\text{ClO}^- / \text{Cl}^-$ electrode. The concentration of $\text{Fe}^{3+}$ is $1000$ times higher than the concentration of $\text{Fe}^{2+}$ in the $\text{Fe}^{3+}/\text{Fe}^{2+}$ electrode. All other conditions remain standard. Use the Nernst equation to calculate the $E$ value of the $\text{Fe}^{3+}/\text{Fe}^{2+}$ electrode. Show your working.
Write an equation for the reaction that occurs in the cell, under these conditions.
A second electrochemical cell is assembled with an $\text{Fe}^{2+}/\text{Fe}$ electrode and an alkaline $\text{ClO}^- / \text{Cl}^-$ electrode under standard conditions. Calculate the value of $\Delta G^\circ$ for the cell.
An iron(II) sulfate solution, $\text{FeSO}_4(aq)$, is electrolysed using iron electrodes. Under the conditions used, no gas is released at the cathode. A current of $0.640\,\text{A}$ is passed for $17.0\,\text{minutes}$. The mass of the cathode rises by $0.185\,\text{g}$. Use these results to calculate an experimental value for the Avogadro constant, $L$. Show your working.
Iron(II) chloride, $\text{FeCl}_2(s)$, is oxidised by chlorine to produce iron(III) chloride, $\text{FeCl}_3(s)$, under standard conditions. $2\text{FeCl}_2(s) + \text{Cl}_2(g) \rightarrow 2\text{FeCl}_3(s)$, $\Delta H^\circ = -128\,\text{kJ mol}^{-1}$. Use Table 3.2 and the other data to calculate the Gibbs free energy change, $\Delta G^\circ$, for this reaction. Show your working.
Iron(II) chloride, $\text{FeCl}_2$, is oxidised by chlorine to produce iron(III) chloride, $\text{FeCl}_3$, under standard conditions. $2\text{FeCl}_2(s)+\text{Cl}_2(g) \rightarrow 2\text{FeCl}_3(s)$ $\Delta H^\circ = -128\,\text{kJ mol}^{-1}$
Predict whether this reaction becomes more or less feasible at a higher temperature. Explain your answer.