What determines if a reaction is spontaneous according to Gibb's Free Energy?

A reaction is considered spontaneous if it has a negative change in Gibbs Free Energy (ΔG). This concept is rooted in thermodynamics, where Gibbs Free Energy helps predict the direction of a reaction under constant temperature and pressure. When ΔG is negative, it indicates that the reaction can occur naturally without the need for external energy input; the system will proceed toward a more stable state.

The relationship between Gibbs Free Energy, enthalpy, and entropy can be summarized by the equation: ΔG = ΔH - TΔS, where ΔH is the change in enthalpy, T is the temperature in Kelvin, and ΔS is the change in entropy. A negative ΔG can result from a reaction that either releases energy (negative ΔH) or increases disorder (positive ΔS), or a combination of both. Thus, the key determinant of spontaneity in this context is that a negative ΔG leads to a thermodynamically favorable reaction that will occur spontaneously.

In contrast, a positive ΔG indicates that the reaction is non-spontaneous and will not occur without an input of energy. The statement regarding enthalpy being greater than entropy does not provide adequate information to determine spontaneity, as both factors need to be considered alongside