The following reaction takes place at 120°C: H₂O(l)$\to$ H₂O(g)$\Delta$H = 44.0 kJ/mol $\Delta$S = 0.119 kJ/mol K Which of the following must be true?

Question

Quizplus · Accepted Answer

The reaction's spontaneity can be determined using the Gibbs free energy change ($\Delta G$), which is calculated using the equation $\Delta G = \Delta H - T\Delta S$. Given that $\Delta H = 44.0 \,kJ/mol$ and $\Delta S = 0.119 \,kJ/mol\,K$, and the temperature is 120°C (or 393K), substituting these values into the equation will show whether $\Delta G$ is positive or negative, indicating spontaneity. However, without doing the calculation, we cannot directly conclude if the reaction is spontaneous ($\Delta G < 0$) or not ($\Delta G > 0$). The information given does not directly confirm any of the statements A, B, C, or D without calculation. However, the question implies a calculation or understanding of thermodynamics is needed, suggesting that more than one of the conceptual statements could be considered correct based on the calculation.

The Following Reaction Takes Place at 120°C: H2O(l) →\to→ H2O(g) Δ\DeltaΔ H = 440 KJ/mol

The Following Reaction Takes Place at 120°C: H₂O(l) $\to$ H₂O(g) $\Delta$ H = 440 KJ/mol