How energy changes decide whether reactions happen
Chemical thermodynamics studies the energy changes in reactions. A system exchanges heat q and work w with its surroundings, and the first law states ΔU = q + w. Most reactions occur at constant pressure, so chemists use enthalpy, H = U + PV, with ΔH = ΔU + ΔnRT for gaseous reactions. A negative ΔH means the reaction is exothermic and releases heat, while a positive ΔH means it is endothermic.
Hess’s law states that the total enthalpy change of a reaction is the same whichever route is taken, so enthalpies of formation, combustion and bond energies can be combined to find unknown values. Spontaneity depends on both enthalpy and entropy, S, a measure of disorder. The Gibbs energy change, ΔG = ΔH − TΔS, decides the outcome: a reaction is spontaneous when ΔG is negative. At equilibrium, ΔG° = −RT ln K, linking thermodynamics to equilibrium constants. For instance, ice melting is endothermic but spontaneous above 0 °C because of the entropy increase.
Why it matters
Thermodynamics tells engineers how much heat a fuel will give, helps food scientists calculate calories, and lets chemists predict whether a reaction is worth attempting at a given temperature. It explains why instant cold packs feel cold and why some reactions need heating. This topic connects closely with equilibrium and electrochemistry.
Key terms
- Enthalpy
- The heat content of a system at constant pressure, H = U + PV.
- Hess’s law
- The total enthalpy change of a reaction is independent of the path taken.
- Entropy
- A measure of the disorder or randomness of a system.
- Gibbs energy
- A function, G = H − TS, whose change decides whether a process is spontaneous.