What occurs when ΔH and ΔS have opposite signs?

When the change in enthalpy (ΔH) and the change in entropy (ΔS) have opposite signs, it significantly impacts the spontaneity of a reaction, as determined by the Gibbs free energy equation, ( \Delta G = \Delta H - T\Delta S ).

If ΔH is positive (endothermic) and ΔS is negative (decrease in disorder), the reaction will not be spontaneous at any temperature. Conversely, if ΔH is negative (exothermic) and ΔS is positive (increase in disorder), the reaction will be spontaneous at all temperatures.

The Gibbs free energy change (ΔG) indicates whether a reaction is spontaneous: it is negative for spontaneous processes. When ΔH and ΔS have opposite signs, it suggests that the reaction's thermodynamic favorability can swing with temperature changes, leading to situations where the reaction is spontaneous at certain temperatures but not others. Therefore, the correct understanding is that within this context, the reaction exhibits temperature-dependent spontaneity.

This relationship showcases a delicate balance between the enthalpic and entropic contributions to the spontaneity of a reaction, and it highlights that these factors must be analyzed collectively to discern the thermodynamic pathway.