If δh°rxn and δs°rxn are both positive values, what drives the spontaneous (favored) reaction and in what direction at standard conditions?

In thermodynamics and physical chemistry, the Gibb's free energy is the criterion for spontaneity. If the Gibb's free energy, denoted as ΔG, is negative, then the reaction is spontaneous. If it is positive, it is non-spontaneous. To estimate ΔG, there is a derived relationship between Gibb's free energy, enthalpy and entropy. The equation is

ΔG=ΔH-TΔS, where ΔH and ΔS is the enthalpy and entropy of the reaction, and T is the temperature at which the reaction happens.
So if ΔH and ΔS are both positive, in order to make ΔG negative, T must be very high. In this way, TΔS would be larger than ΔH.

Therefore, a reaction with a positive ΔH and ΔS will be spontaneous at high temperatures.

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