ExcelleThermodynamics explains why and how matter changes from one phase to another (solid, liquid, gas, plasma) by looking at energy, enthalpy, entropy, and Gibbs free energy. Here’s the breakdown:
1. Energy and Enthalpy in Phase Changes
- Phase changes involve heat transfer without a temperature change (at constant pressure).
- Enthalpy (ΔH) quantifies the heat absorbed or released:
- Fusion (melting) → requires heat input (ΔH > 0).
- Vaporization (boiling/evaporation) → requires more heat input (ΔH > 0).
- Condensation and freezing → release heat (ΔH < 0).
2. Entropy (ΔS) and Disorder
- Entropy measures randomness or molecular freedom.
- During phase change:
- Solid → Liquid → Gas → Entropy increases (more disorder).
- Gas → Liquid → Solid → Entropy decreases (more order).
3. Gibbs Free Energy (ΔG) and Spontaneity
- The equation ΔG = ΔH – TΔS determines if a phase change happens spontaneously at a given temperature and pressure.
- At equilibrium (like melting point or boiling point): ΔG = 0.
- Example: Ice at 0 °C can coexist with water because ΔG = 0 for melting/freezing.
4. Phase Equilibria and Diagrams
- Thermodynamics gives phase diagrams showing conditions (temperature, pressure) where phases are stable.
- Example: Water boils at 100 °C at 1 atm because vapor and liquid phases have the same Gibbs free energy there.
5. Latent Heat
- Heat absorbed or released during phase change is called latent heat, and it comes directly from enthalpy change (ΔH).
- Heat of fusion (melting/freezing).
- Heat of vaporization (boiling/condensation).
In short:
- Enthalpy explains heat absorbed/released.
- Entropy explains disorder change.
- Gibbs free energy predicts when the phase change occurs.