Final answer:
The reaction is non-spontaneous under standard conditions because it has a positive ΔH and a negative ΔS, resulting in a positive Gibbs free energy change. Spontaneity may vary with temperature.
Step-by-step explanation:
When evaluating the spontaneity of a chemical reaction, the enthalpy change (ΔH) and entropy change (ΔS) play significant roles. In this case, the reaction has an enthalpy change of ΔH = 620 kJ/mol, which is a positive value, indicating that the reaction is endothermic and requires heat from the surroundings. The entropy change for the reaction is ΔS = -0.46 kJ/(mol*K), which is negative, suggesting a decrease in disorder or randomness of the system as the reaction proceeds.
According to the Gibbs free energy equation (ΔG = ΔH - TΔS), the sign of the free energy change (ΔG) determines whether a reaction is spontaneous. At constant temperature and pressure, a reaction is spontaneous if ΔG is negative. Since ΔH is positive and ΔS is negative, the product of TΔS will also be negative, leading to a positive ΔG. This suggests that the reaction is non-spontaneous under standard conditions. However, it's important to note that the spontaneity might change at different temperatures, as the temperature (T) can influence the TΔS term and potentially make ΔG negative.