a) i) Define bond energy. Bond energy is the average enthalpy change required to break one mole of a specific type of covalent bond in a gaseous molecule. a) ii) Write an equation to represent the lattice energy of calcium fluoride. $$ \text{Ca}^{2+}(\text{g}) + 2\text{F}^{-}(\text{g}) \longrightarrow \text{CaF}_{2}(\text{s}) $$ a) iii) Calculate the total heat of solution of sodium chloride and hence construct an energy cycle for the dissolution. Step 1: State the given enthalpy values. Lattice energy of NaCl, $\Delta H_{L} = -783 \text{ kJ/mol}$ (This is the enthalpy change for forming NaCl(s) from gaseous ions). Enthalpy of hydration of Na$^{+}$, $\Delta H_{hyd} [\text{Na}^{+}(\text{g})] = -402 \text{ kJ/mol}$ Enthalpy of hydration of Cl$^{-}$, $\Delta H_{hyd} [\text{Cl}^{-}(\text{g})] = -377 \text{ kJ/mol}$ Step 2: Apply Hess's Law to calculate the enthalpy of solution ($\Delta H_{soln}$). The dissolution process is: $\text{NaCl(s)} \longrightarrow \text{Na}^{+}(\text{aq}) + \text{Cl}^{-}(\text{aq})$ This can be broken down into two steps: 1. Breaking the lattice into gaseous ions: $\text{NaCl(s)} \longrightarrow \text{Na}^{+}(\text{g}) + \text{Cl}^{-}(\text{g})$. The enthalpy change for this step is the negative of the lattice energy, $-\Delta H_{L}$. 2. Hydrating the gaseous ions: $\text{Na}^{+}(\text{g}) + \text