enthalpy–entropy balance
Whether a change happens on its own is settled by a tug-of-war between two pulls. One is enthalpy: nature likes to release heat and slide to lower energy, the way a ball rolls into a dip. The other is entropy: nature likes to spread out, multiply its options, and grow more disordered. Sometimes both pull the same way; often they pull against each other, and the winner depends on the temperature.
The scoreboard is the Gibbs equation, ΔG = ΔH − TΔS. The enthalpy term ΔH counts the energy released or absorbed; the entropy term TΔS weighs the change in disorder, scaled by the absolute temperature T. A change is spontaneous only when the combination comes out negative. Because T multiplies the entropy term, temperature acts as the referee: it decides how much weight entropy gets in the contest.
Why it matters: this single balance explains why some reactions reverse direction as you heat them. When ΔH and ΔS have the same sign, raising the temperature flips which term dominates — so ice melts above 0 °C (entropy wins) but freezes below it (enthalpy wins). The honest nuance: the entropy that matters is the total for the universe, but folding the surroundings' share into −TΔS lets you score the whole contest using the system alone.
Dissolving ammonium nitrate soaks up heat (ΔH positive, unfavorable) yet still happens, because the salt spreading through water raises entropy so much that TΔS overpowers ΔH — and the beaker turns cold.
Entropy can win even when enthalpy says no.
A quick map of the four cases: if ΔH is negative and ΔS positive, the change is spontaneous at all temperatures; if both are positive, only when hot; if both negative, only when cold; and if ΔH is positive with ΔS negative, never on its own.