Final answer:
For ∆G° to be negative at all temperatures, ∆H° must be negative (exothermic) and ∆S° must be positive (entropy increasing). This results in a reaction that is spontaneous in the forward direction at any temperature.
Step-by-step explanation:
If ∆G° < 0 for a reaction at all temperatures, then we can conclude that ∆H° is negative and ∆S° is positive. This is because the Gibbs free energy equation, which is ∆G° = ∆H° - T∆S°, dictates that for ∆G° to be negative at all temperatures, the enthalpy (∆H°) must be releasing heat (exothermic), which is indicated by a negative value. Conversely, the entropy (∆S°) must be increasing, as shown by a positive value, for the -T∆S° term to be sufficiently negative to overcome any positive enthalpy change and still result in a negative ∆G°. As temperature in this equation is in Kelvin and cannot be negative, a negative ∆H° and a positive ∆S° ensure that the reaction is spontaneously forward at any temperature.