What can you infer if both ΔH and ΔS are positive for a reaction?

When both the change in enthalpy (ΔH) and the change in entropy (ΔS) for a reaction are positive, we can analyze the spontaneity of the reaction using the Gibbs free energy equation:

[ ΔG = ΔH - TΔS ]

Here, ΔG represents the change in Gibbs free energy, T is the temperature in Kelvin, and the signs of ΔH and ΔS will influence whether ΔG is positive or negative. For a reaction to be spontaneous, ΔG must be negative.

With both ΔH and ΔS being positive, at low temperatures, the term (TΔS) can be small, which means ΔH will dominate the equation. This can potentially result in a positive ΔG, indicating non-spontaneity. However, as the temperature increases, the (TΔS) term becomes larger and will eventually outweigh ΔH. This leads to a negative ΔG, indicating that the reaction will become spontaneous as temperature increases.

Consequently, as temperature rises, the likelihood of spontaneity increases, which aligns with the correct answer. The subtle balance driven by temperature shows how favorable entropy changes can ultimately drive a reaction toward spontaneity, even when both ΔH and ΔS