Introduction
Chemical thermodynamics asks whether a reaction can happen at all (spontaneity) and how much energy it releases or absorbs, without needing to know how fast it proceeds (that is kinetics, Chapter 4). It uses the same laws met in Physics Chapter IV, applied specifically to chemical systems.
1. System, surroundings, and state functions
A system is the part of the universe under study; everything else is the surroundings. A state function (like internal energy, enthalpy, or entropy) depends only on the current state of the system, not on the path taken to reach it — this is what makes Hess’s law (below) possible: you can add up steps in any convenient path because the total change only depends on start and end points.
2. The laws of thermodynamics in chemical systems
The first law (ΔU = Q − W) means the energy change of a reaction must appear as heat exchanged with the surroundings plus any work done (commonly gas expansion/compression). The second law states that the total entropy of the universe increases for any spontaneous process — chemical reactions proceed in the direction that increases overall (system + surroundings) disorder. The third law sets absolute zero entropy for a perfect crystal at 0 K, providing the reference point needed to calculate absolute entropies.
3. Internal energy, enthalpy, entropy, and Gibbs free energy
Enthalpy (H = U + PV) is convenient because, at constant pressure (typical for reactions open to the atmosphere), the enthalpy change equals the heat exchanged — exothermic reactions have negative ΔH, endothermic positive. Entropy (S) measures disorder/dispersal of energy. Gibbs free energy (G = H − TS) combines both into a single spontaneity criterion: ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG < 0, regardless of whether it is exothermic or endothermic, because a sufficiently large favourable entropy change can drive an endothermic reaction forward (and vice versa).
4. Heat capacities and thermochemical relations
Heat capacity is the heat required to raise a system’s temperature by one degree; it differs at constant volume (Cv) versus constant pressure (Cp) because constant-pressure heating also does expansion work. Hess’s law — a direct consequence of enthalpy being a state function — allows the enthalpy of a reaction to be calculated by adding together the enthalpies of a convenient sequence of steps that achieve the same overall transformation, even if that sequence is not how the reaction actually occurs.
5. Criteria for spontaneity and equilibrium
The sign of ΔG determines spontaneity: negative means spontaneous as written, positive means the reverse reaction is spontaneous, and ΔG = 0 defines equilibrium — the point at which forward and reverse reaction tendencies exactly balance, connecting thermodynamics directly to the equilibrium constant, K, via ΔG° = −RT ln K.
Key Takeaways
- State functions (U, H, S, G) depend only on current state, enabling path-independent calculations like Hess’s law.
- The first law tracks energy conservation; the second law, via entropy, determines the direction spontaneous processes take.
- Gibbs free energy, ΔG = ΔH − TΔS, is the single criterion for spontaneity, showing that both enthalpy and entropy changes matter, not enthalpy alone.
- ΔG = 0 defines equilibrium and links thermodynamics quantitatively to the equilibrium constant K.