This question is about energy changes. Nitrogen monoxide, NO\text{NO}NO, decomposes as shown in Reaction 1:
NO(g)→12N2(g)+12O2(g)Reaction 1 \text{NO}(\text{g}) \rightarrow \frac{1}{2}\text{N}_2(\text{g}) + \frac{1}{2}\text{O}_2(\text{g}) \quad \text{\textbf{Reaction 1}} NO(g)→21N2(g)+21O2(g)Reaction 1(a) The table shows enthalpy changes of formation and standard entropies:
| Substance | ΔHf⊖\Delta H_{\text{f}}^\ominusΔHf⊖ / kJ mol−1\text{kJ mol}^{-1}kJ mol−1 | S⊖S^\ominusS⊖ / J K−1 mol−1\text{J K}^{-1}\text{ mol}^{-1}J K−1 mol−1 |
|---|---|---|
| NO(g)\text{NO}(\text{g})NO(g) | +90.4+90.4+90.4 | 211211211 |
| N2(g)\text{N}_2(\text{g})N2(g) | 000 | 192192192 |
| O2(g)\text{O}_2(\text{g})O2(g) | 000 | 205205205 |
Calculate the free-energy change, ΔG\Delta GΔG, in kJ mol−1\text{kJ mol}^{-1}kJ mol−1, of Reaction 1 at 25 ∘C25\ ^\circ\text{C}25 ∘C. Give your answer to 3 significant figures.
The decomposition of nitrogen monoxide shown in Reaction 1 is thermodynamically feasible. Suggest why Reaction 1 does not take place at 25 ∘C25\ ^\circ\text{C}25 ∘C despite being feasible.
