ABSTRACT
At thermodynamic equilibrium, a material predominately exists in the physical state (gas, liquid, solid) with the lowest free energy (McClements 2012). The free energy of a particular state is governed by enthalpy (energy stored in molecular interactions) and entropy (disorder) contributions. The free energy change (ΔGT) when a system undergoes a transition from one state to another can be conveniently represented by the following expression:
ΔGT = ΔHT − T × ΔST. (10.1)
Here, ΔHT is the change in enthalpy (molecular interaction strength) associated with the transition and ΔST is the change in entropy (disorder) associated with the transition. A phase transition is thermodynamically favorable when ΔGT < 0 but is thermodynamically unfavorable when ΔGT > 0. At the transition temperature (Tm), ΔGT = 0, and each state is equally favorable. The amount of energy stored in the molecular interactions in different states typically follows the order: solids > liquids > gas, whereas the entropy of different states typically follows the order: gas > liquids > solids. Thus, at relatively low temperatures (small T ), the enthalpy contribution dominates (since strong attractive interactions means a more negative ΔHT), which favors the existence of the solid phase. Conversely, at high temperatures (large T ), the entropy contribution dominates (−T × ΔST), which favors the existence of the gas phase. At intermediate temperatures, the balance of enthalpy and entropy effects favors the existence of the liquid phase.
