Final answer:
The spontaneity of a reaction depends on the sign and magnitude of enthalpy (ΔH) and entropy (ΔS) changes, and the temperature. High temperatures favor reactions with positive ΔS, while low temperatures favor reactions with negative ΔH. The standard free energy change (ΔG) relates to equilibrium constants, with negative ΔG corresponding to an equilibrium constant greater than 1.
Step-by-step explanation:
Temperature plays a significant role in determining the spontaneity of a chemical reaction, according to the Gibbs free energy equation ΔG = ΔH - TΔS, where ΔG is the change in free energy, ΔH is the enthalpy change, T is the temperature in Kelvin, and ΔS is the entropy change. At relatively high temperatures, a reaction with a positive ΔH (endothermic reaction) and a large positive ΔS (increase in entropy) is more likely to be spontaneous because the TΔS term will have a great enough positive value to outweigh the positive ΔH, making ΔG negative which indicates spontaneity. Conversely, at relatively low temperatures, a reaction with a negative ΔH (exothermic reaction) and a positive ΔS is likely to be spontaneous because the negative ΔH will dominate the ΔG term even when the TΔS term is small. This relationship also relates the standard free energy changes to equilibrium constants, where a negative ΔG corresponds to a value of the equilibrium constant greater than 1, indicating a spontaneous reaction under standard conditions.