Chemical Energetics

Define enthalpy change of hydration.

Explain the relative magnitudes of the enthalpy changes of hydration of $\text{K}^+$ and $\text{Ca}^{2+}$.

Define lattice energy.

The lattice energy, $\Delta H_{\text{latt}}$, of calcium fluoride, $\text{CaF}_2$, is $-2602 \text{ kJ mol}^{-1}$.

Calculate the enthalpy change of solution, $\Delta H_{\text{sol}}$, in $\text{kJ mol}^{-1}$, of $\text{CaF}_2$.

The formation of $\text{CaF}_2$ at 298 K is shown.

Calculate the entropy change, $\Delta S^\ominus$, in $\text{J K}^{-1} \text{ mol}^{-1}$, for this reaction.

Define entropy.

Identify the process occurring at each of the temperatures $T_1$ and $T_2$.

$T_1$ .................................................. $T_2$ ..................................................

Explain why the entropy change, $\Delta S$, at $T_2$ is bigger than the entropy change at $T_1$.

The equation for the reduction of iron(III) oxide by carbon monoxide at $450\text{ }^\circ\text{C}$ is shown.

Table 3.1 shows the enthalpy of formation, $\Delta H_f^\ominus$, and the entropy, $S^\ominus$, for some substances.

Table 3.1

Use the data in Table 3.1 to calculate the entropy, $S^\ominus$, of carbon monoxide at $450\text{ }^\circ\text{C}$.

Show your working.

Iron(II) oxide can also be reduced to iron by carbon monoxide, as shown.

State the effect of increasing temperature on the feasibility of this reaction.

Explain your answer.

Define enthalpy change of hydration.

Explain the relative magnitudes of the enthalpy changes of hydration of $\text{K}^+$ and $\text{Ca}^{2+}$.

Define lattice energy.

The lattice energy, $\Delta H_{\text{latt}}$, of calcium fluoride, $\text{CaF}_2$, is $-2602\text{ kJ mol}^{-1}$.

Calculate the enthalpy change of solution, $\Delta H_{\text{sol}}$, in $\text{kJ mol}^{-1}$, of $\text{CaF}_2$.

$\Delta H_{\text{sol}}$ of $\text{CaF}_2$ = .............................. $\text{kJ mol}^{-1}$

The formation of $\text{CaF}_2$ at 298 K is shown.

Calculate the entropy change, $\Delta S^\ominus$, in $\text{J K}^{-1}\text{ mol}^{-1}$, for this reaction.

$\Delta S^\ominus$ = .............................. $\text{J K}^{-1}\text{ mol}^{-1}$

Define lattice energy, $\Delta H_\text{latt}$.

Define enthalpy change of solution, $\Delta H_\text{sol}$.

The enthalpy change of hydration can be represented by $\Delta H_\text{hyd}$.

Write the mathematical expression for the $\Delta H_\text{sol}$ of $\text{NaCl}$ in terms of $\Delta H_\text{latt}(\text{NaCl})$, $\Delta H_\text{hyd}(\text{Na}^+)$ and $\Delta H_\text{hyd}(\text{Cl}^-)$.

Complete the Born–Haber cycle in Fig. 8.1 for the ionic solid $\text{NaCl}$.

Include state symbols of relevant species.

Predict which of the ions, $\text{Cl}^-$ or $\text{NO}_3^-$, has the more negative enthalpy change of hydration.

Explain your answer.

Define enthalpy change of atomisation, $\Delta H_{at}$.

Define first electron affinity, EA.

Explain why the first electron affinity of chlorine is more exothermic than the first electron affinity of iodine.

The enthalpy change for the reaction $\text{Cl}_2\text{(g)} + 2\text{e}^- \rightarrow 2\text{Cl}^-\text{(g)}$ is $-486 \text{ kJ mol}^{-1}$.

The first electron affinity of chlorine is $-364 \text{ kJ mol}^{-1}$.

Calculate the enthalpy change of atomisation of chlorine.

$\Delta H_{at}$ of chlorine = .............................. $\text{kJ mol}^{-1}$

Define the term entropy.

Place one tick ($\checkmark$) in each row of Table 3.1 to show the sign of the entropy change, $\Delta S$, for each process.

Table 3.1

Chlorine trifluoride, $\text{ClF}_3$, decomposes on heating into its elements, as shown.

Standard entropies are shown in Table 3.2.

Table 3.2

Calculate the standard entropy change, $\Delta S^\ominus$, in $\text{J K}^{-1} \text{mol}^{-1}$, for reaction 1.

Predict the sign of the entropy change, $\Delta S$, for reaction 2.

Explain your answer.

The Gibbs equation is shown.

Fig. 3.1 shows values of the Gibbs free energy change, $\Delta G^\ominus$, in $\text{kJ mol}^{-1}$, at different temperatures, $T$, in $\text{K}$, for reaction 2.

Assume $\Delta H^\ominus$ and $\Delta S^\ominus$ values for this reaction remain constant over this temperature range.

Use the gradient and intercept on the $y$-axis in Fig. 3.1 and the Gibbs equation to determine:

• $\Delta S^\ominus$, in $\text{J K}^{-1} \text{mol}^{-1}$, for reaction 2
• the minimum temperature, $T$, in $\text{K}$, at which the reaction is feasible
• $\Delta H^\ominus$, in $\text{kJ mol}^{-1}$, for reaction 2.

Carbon disulfide, $\text{CS}_2$, is flammable and reacts readily with oxygen, as shown in reaction 1.

Table 3.1 shows the standard enthalpy of formation, $\Delta H_{\text{f}}^{\ominus}$, and the standard entropy, $S^{\ominus}$, for some substances.

Table 3.1

Calculate the standard Gibbs free energy change, $\Delta G^{\ominus}$, in $\text{kJ mol}^{-1}$, for reaction 1 at $25\text{ }^{\circ}\text{C}$.

Carbon disulfide reacts with chlorine to form tetrachloromethane, as shown in reaction 2.

Calculate the maximum temperature, in $\text{K}$, for reaction 2 to be feasible.

Predict and explain the variation in enthalpy change of hydration for the ions $\text{F}^-$, $\text{Cl}^-$, $\text{Br}^-$ and $\text{I}^-$.

Complete line D. Include state symbols.

The value of the enthalpy change for process 1 can be calculated using the values of five other enthalpy changes which are not referred to in Fig. 2.1.

Identify these five other enthalpy changes, using either names or symbols.

Define lattice energy, $\Delta H_{\text{latt}}$.

Complete the expression to give the mathematical relationship between $\Delta H_{\text{latt}}$ of calcium fluoride and the enthalpy changes for processes 1 and 3.

Use data from Table 2.1 to calculate a value for the hydration energy, $\Delta H_{\text{hyd}}$, of fluoride ions, $\text{F}^-\text{(g)}$.

Table 2.1

Define entropy.

At $298\text{ K}$, the Gibbs free energy change, $\Delta G$, for the solution of compound $\mathbf{T}$ is $+6.00\text{ kJ}\,\text{mol}^{-1}$.

The enthalpy change of solution, $\Delta H_{\text{sol}}$, of compound $\mathbf{T}$ is $+30.0\text{ kJ}\,\text{mol}^{-1}$ at $298\text{ K}$.

Calculate the value of the entropy change, $\Delta S$, for the solution of compound $\mathbf{T}$ at $298\text{ K}$.

Predict whether compound $\mathbf{T}$ becomes more or less soluble as the water is heated from $298\text{ K}$ to $360\text{ K}$. Explain your answer.

Predict and explain the variation in enthalpy change of hydration for the ions $\text{Na}^+$, $\text{Mg}^{2+}$ and $\text{Al}^{3+}$.

Complete line C on Fig. 2.1. Include state symbols.

Use both words and symbols to identify change 2 on Fig. 2.1.

Use changes 1 and 3 as examples of how this should be done.

Calculate a value for the lattice energy of magnesium chloride, $\Delta H_{\text{latt}}\text{ MgCl}_2(\text{s})$, by selecting and using appropriate data from Table 2.1.

Table 2.1

Define entropy.

At $25^\circ\text{C}$ the enthalpy change of solution of compound $\mathbf{Z}$ is $+26\text{ kJ mol}^{-1}$. The entropy change of solution of $\mathbf{Z}$ at the same temperature is $+52\text{ J K}^{-1}\text{mol}^{-1}$.

Calculate the value of the Gibbs free energy change, $\Delta G$, for the solution of $\mathbf{Z}$ at $25^\circ\text{C}$.

Use your answer to (d) to predict whether or not $\mathbf{Z}$ is soluble in water at $25^\circ\text{C}$. Explain your answer.

Predict whether $\mathbf{Z}$ becomes more or less soluble as the water is heated from $25^\circ\text{C}$ to $95^\circ\text{C}$. Explain your answer.

Predict and explain the variation in enthalpy change of hydration for the ions $\text{F}^-$, $\text{Cl}^-$, $\text{Br}^-$ and $\text{I}^-$.

Complete line D. Include state symbols.

The value of the enthalpy change for process 1 can be calculated using the values of five other enthalpy changes which are not referred to in Fig. 2.1.

Identify these five other enthalpy changes, using either names or symbols.

Define lattice energy, $\Delta H_{\text{latt}}$.

Complete the expression to give the mathematical relationship between $\Delta H_{\text{latt}}$ of calcium fluoride and the enthalpy changes for processes 1 and 3.

Use data from Table 2.1 to calculate a value for the hydration energy, $\Delta H_{\text{hyd}}$, of fluoride ions, $\text{F}^-\text{(g)}$.

Table 2.1

Define entropy.

At $298\text{ K}$, the Gibbs free energy change, $\Delta G$, for the solution of compound T is $+6.00\text{ kJ mol}^{-1}$.

The enthalpy change of solution, $\Delta H_{\text{sol}}$, of compound T is $+30.0\text{ kJ mol}^{-1}$ at $298\text{ K}$.

Calculate the value of the entropy change, $\Delta S$, for the solution of compound T at $298\text{ K}$.

Predict whether compound T becomes more or less soluble as the water is heated from $298\text{ K}$ to $360\text{ K}$. Explain your answer.