Define entropy in terms of possible arrangements and energy distribution.
Identify the phase change occurring at each of the temperatures $T_1$ and $T_2$.
Explain why the entropy change, $\Delta S$, at $T_2$ is greater than the entropy change at $T_1$.
The equation for the reduction of iron(III) oxide by carbon monoxide at $450\,^{\circ}\text{C}$ is shown. $\text{Fe}_2\text{O}_3(s) + 3\text{CO}(g) \rightarrow 2\text{Fe}(s) + 3\text{CO}_2(g)$ \quad $\Delta G^\circ = -36.2\,\text{kJ mol}^{-1}$ Table 3.1 lists the enthalpy of formation, $\Delta H_f^\circ$, and the entropy, $S^\circ$, for some substances. Use the information in Table 3.1 to calculate the entropy, $S^\circ$, of carbon monoxide at $450\,^{\circ}\text{C}$. Show your working clearly.
Iron(II) oxide can also be reduced by carbon monoxide, as shown. $\text{FeO}(s) + \text{CO}(g) \rightarrow \text{Fe}(s) + \text{CO}_2(g)$ $\Delta H^\circ = -11.1\,\text{kJ mol}^{-1}$ $\Delta S^\circ = -15.2\,\text{J K}^{-1}\text{ mol}^{-1}$ State the effect of increasing temperature on the feasibility of this reaction. Explain your answer.