Question #44128

Chromium(VI) oxide (\ce{CrO3}) is a powerful oxidizing agent that undergoes thermal decomposition to chromium(III) oxide (\ce{Cr2O3}) and oxygen gas. Consider the balanced chemical equation for the decomposition of 2 moles of \ce{CrO3(s)}: $\ce{2CrO3(s) -> Cr2O3(s) + 3/2 O2(g)}$ Given the following thermodynamic data at $298 \text{ K}$: $\Delta H_f^\circ(\ce{CrO3}, s) = -590 \text{ kJ mol}^{-1}$ $\Delta H_f^\circ(\ce{Cr2O3}, s) = -1060 \text{ kJ mol}^{-1}$ $S^\circ(\ce{CrO3}, s) = 70 \text{ J K}^{-1}\text{mol}^{-1}$ $S^\circ(\ce{Cr2O3}, s) = 65 \text{ J K}^{-1}\text{mol}^{-1}$ $S^\circ(\ce{O2}, g) = 210 \text{ J K}^{-1}\text{mol}^{-1}$ Assuming $\Delta H^\circ$ and $\Delta S^\circ$ are independent of temperature, calculate the standard Gibbs free energy change ($\Delta G^\circ$) for this reaction at $600 \text{ K}$ and identify the primary thermodynamic driving force for the decomposition.