A sample of hydrated lithium ethanedioate, $\text{Li}_2\text{C}_2\text{O}_4\cdot\text{H}_2\text{O}$, is gently heated, producing two gases and leaving a white solid behind. The solid is then added to $\text{HNO}_3\text{(aq)}$. A gas is evolved that turns limewater milky. Complete the equation for the decomposition of $\text{Li}_2\text{C}_2\text{O}_4\cdot\text{H}_2\text{O}$.
The pattern in decomposition temperatures for the Group 2 ethanedioates is similar to that for the Group 2 nitrates. Suggest which of $\text{CaC}_2\text{O}_4$ and $\text{BaC}_2\text{O}_4$ will decompose at the lower temperature. Explain your answer.
Potassium iron(III) ethanedioate, $\text{K}_3[\text{Fe(C}_2\text{O}_4)_3]$, dissolves in water to give a green solution. Explain why transition elements can form coloured complexes.
When heated, the anhydrous iron(III) compound $\text{K}_3[\text{Fe(C}_2\text{O}_4)_3]$ decomposes to form a mixture of $\text{K}_2[\text{Fe(C}_2\text{O}_4)_2]$, $\text{K}_2\text{C}_2\text{O}_4$ and $\text{CO}_2$. Complete the equation for the decomposition of $\text{K}_3[\text{Fe(C}_2\text{O}_4)_3]$.
The $[\text{Fe}(\text{C}_2\text{O}_4)_3]^{3-}$ complex ion displays stereoisomerism. Complete the three-dimensional diagrams in Fig. 3.1 to show the two stereoisomers of $[\text{Fe}(\text{C}_2\text{O}_4)_3]^{3-}$. The $\text{C}_2\text{O}_4^{2-}$ ligand may be represented using O-O.
Buffer solutions are used to control pH. Write two equations to describe how a solution containing $\text{HC}_2\text{O}_4^-$ ions behaves as a buffer solution when small amounts of acid or alkali are added.
A fuel cell is an electrochemical cell that can be used to generate electrical energy by using oxygen to oxidise a fuel. Ethanedioic acid, $(\text{COOH})_2$, dissolved in an alkaline electrolyte is being investigated as a fuel. The relevant standard electrode potentials, $E^\circ$, for the cell are shown. $\text{O}_2\text{(g)} + 2\text{H}_2\text{O(l)} + 4e^- \rightleftharpoons 4\text{OH}^-\text{(aq)}$, $E^\circ = +0.40\,\text{V}$. $2\text{CO}_2\text{(g)} + 2e^- \rightleftharpoons \text{C}_2\text{O}_4^{2-}\text{(aq)}$, $E^\circ = -0.59\,\text{V}$. Use these equations to deduce the overall cell reaction and calculate the value of $E^\circ_{\text{cell}}$.