How does the spontaneity of a reaction change with temperature?

The spontaneity of a reaction can increase or decrease with temperature, depending on the reaction's enthalpy and entropy changes.

The spontaneity of a chemical reaction is determined by the Gibbs free energy change (ΔG), which is calculated using the equation ΔG = ΔH - TΔS. Here, ΔH is the change in enthalpy, T is the absolute temperature in Kelvin, and ΔS is the change in entropy.

If a reaction has a negative ΔH (exothermic) and positive ΔS (increase in disorder), it will be spontaneous at all temperatures. This is because both terms in the equation contribute to a negative ΔG, which indicates spontaneity.

However, if a reaction has a positive ΔH (endothermic) and negative ΔS (decrease in disorder), it will be non-spontaneous at all temperatures. Both terms in the equation contribute to a positive ΔG, indicating non-spontaneity.

The interesting cases are when ΔH and ΔS have opposite signs. If ΔH is positive and ΔS is negative, the reaction is spontaneous at low temperatures. This is because the TΔS term is small at low temperatures, so the positive ΔH dominates, leading to a negative ΔG. Conversely, if ΔH is negative and ΔS is positive, the reaction is spontaneous at high temperatures. Here, the TΔS term becomes large at high temperatures, so the negative ΔH dominates, leading to a negative ΔG.

In summary, the spontaneity of a reaction can change with temperature, and this is determined by the relative magnitudes and signs of the enthalpy and entropy changes. Understanding this concept is crucial for predicting whether a reaction will occur under certain conditions, and it is a fundamental aspect of thermodynamics in chemistry.