The spontaneity of a chemical reaction can vary with temperature, influenced by the changes in enthalpy and entropy associated with the reaction.

The spontaneity of a reaction is quantified by the Gibbs free energy change, denoted as $\Delta G$. The relationship is expressed by the equation:

In this equation, $\Delta H$ represents the change in enthalpy, $T$ is the absolute temperature measured in Kelvin, and $\Delta S$ denotes the change in entropy.

For a reaction that has a negative $\Delta H$ (indicating it is exothermic) and a positive $\Delta S$ (indicating an increase in disorder), the reaction will be spontaneous at all temperatures. This is because both terms in the equation contribute to a negative $\Delta G$, which signifies spontaneity.

Conversely, if a reaction exhibits a positive $\Delta H$ (indicating it is endothermic) and a negative $\Delta S$ (indicating a decrease in disorder), the reaction will be non-spontaneous at all temperatures. In this case, both terms lead to a positive $\Delta G$, indicating non-spontaneity.

The more intriguing scenarios arise when $\Delta H$ and $\Delta S$ have opposite signs. If $\Delta H$ is positive and $\Delta S$ is negative, the reaction will be spontaneous at low temperatures. This occurs because the term $T \Delta S$ is small at low temperatures, allowing the positive $\Delta H$ to dominate, resulting in a negative $\Delta G$. Conversely, if $\Delta H$ is negative and $\Delta S$ is positive, the reaction is spontaneous at high temperatures. In this situation, the term $T \Delta S$ becomes significant at elevated temperatures, allowing the negative $\Delta H$ to dominate, thus leading to a negative $\Delta G$.

In summary, the spontaneity of a reaction can indeed change with temperature, which is determined by the relative magnitudes and signs of the changes in enthalpy and entropy. Grasping this concept is essential for predicting whether a reaction will occur under specific conditions; it is a fundamental principle of thermodynamics in chemistry.