IB Chemistry – Reactivity 1: What Drives Reactions

In a calorimetry experiment, burning a small sample of fuel raises the temperature of $100\,\text{g}$ of water by $6.5\degree\text{C}$. The mass of fuel burned corresponds to $0.0050\,\text{mol}$.(a) Calculate the heat energy transferred to the water, in kJ.(b) Calculate the molar enthalpy of combustion of the fuel, including its sign.

State Hess’s law, and explain why it is a consequence of the first law of thermodynamics (conservation of energy).

Given the standard enthalpies of combustion $\Delta H_c^{\ominus}[\text{C(s)}]=-393.5\,\text{kJ mol}^{-1}$, $\Delta H_c^{\ominus}[\text{H}_2\text{(g)}]=-285.8\,\text{kJ mol}^{-1}$ and $\Delta H_c^{\ominus}[\text{CH}_4\text{(g)}]=-890.3\,\text{kJ mol}^{-1}$, use Hess’s law to calculate the standard enthalpy of formation of methane, $\Delta H_f^{\ominus}[\text{CH}_4\text{(g)}]$, for the reaction $\text{C(s)}+2\text{H}_2\text{(g)}\rightarrow\text{CH}_4\text{(g)}$.

Explain, in terms of bond breaking and bond making, why some reactions are exothermic and others are endothermic.

Use the following data to construct a Born-Haber cycle for sodium chloride and calculate the lattice enthalpy of NaCl:Enthalpy of atomisation of Na(s): $+107\,\text{kJ mol}^{-1}$Enthalpy of atomisation of $\frac{1}{2}\text{Cl}_2$(g): $+122\,\text{kJ mol}^{-1}$First ionisation energy of Na(g): $+496\,\text{kJ mol}^{-1}$First electron affinity of Cl(g): $-349\,\text{kJ mol}^{-1}$Enthalpy of formation of NaCl(s): $-411\,\text{kJ mol}^{-1}$

Define lattice enthalpy, and explain why lattice enthalpy values are always exothermic (negative).

For the reaction $\text{N}_2\text{(g)}+3\text{H}_2\text{(g)}\rightleftharpoons2\text{NH}_3\text{(g)}$, $\Delta H^{\ominus}=-92\,\text{kJ mol}^{-1}$ and $\Delta S^{\ominus}=-198\,\text{J K}^{-1}\text{mol}^{-1}$.(a) Calculate $\Delta G^{\ominus}$ at $298\,\text{K}$, and state whether the reaction is spontaneous at this temperature.(b) Calculate the temperature above which the reaction becomes non-spontaneous under standard conditions.

Explain why $\Delta S^{\ominus}$ is negative for the reaction $\text{N}_2\text{(g)}+3\text{H}_2\text{(g)}\rightleftharpoons2\text{NH}_3\text{(g)}$.

A voltaic cell is set up using the half-cells $\text{Zn}^{2+}\text{(aq)}+2e^-\rightleftharpoons\text{Zn(s)}$, $E^{\ominus}=-0.76\,\text{V}$, and $\text{Cu}^{2+}\text{(aq)}+2e^-\rightleftharpoons\text{Cu(s)}$, $E^{\ominus}=+0.34\,\text{V}$.(a) Calculate the standard cell potential, $E_{\text{cell}}^{\ominus}$.(b) State which electrode is the anode (oxidation) and which is the cathode (reduction).

Using $\Delta G^{\ominus}=-nFE_{\text{cell}}^{\ominus}$, calculate the standard Gibbs energy change, in $\text{kJ mol}^{-1}$, for the cell reaction in the previous question ($n=2$).

Explain the relationship between the sign of $E_{\text{cell}}^{\ominus}$ and the sign of $\Delta G^{\ominus}$, and state what this tells you about whether a redox reaction is spontaneous.

The enthalpy of solution of an ionic compound can be found from a Hess’s law cycle linking lattice enthalpy and enthalpy of hydration. Describe this cycle, and explain, in terms of these two enthalpy changes, why some ionic compounds dissolve exothermically while others dissolve endothermically.