Phase Transitions

Whether solid, liquid, or gas, a substance's state depends on the order and arrangement of its particles (atoms, molecules, or ions). Particles in the solid pack closely together, generally in a pattern. The particles vibrate about their fixed positions but do not move or squeeze past their neighbors. In liquids, although the particles are closely spaced, they are randomly arranged. The position of the particles are not fixed—that is, they are free to move past their neighbors to occupy different locations. Because the particles are close together in the solid and liquid states, these are referred to as the condensed states or condensed phases. In these states, substances exhibit relatively strong intermolecular forces. In gases, the interparticle forces of attractions are weak. The particles of a gas are not constrained by their neighbors; the particles are free to move and, under normal conditions, are separated by large distances.

The internal energy of a substance—the total kinetic energy of all its molecules—depends on the strength of the intermolecular forces in the condensed phases and the pressure exerted on the substance. The internal energy of a substance is the highest in a gaseous state, the lowest in a solid state, and intermediate in a liquid.

Phase transitions are caused by changes in physical conditions, such as in temperature and/or pressure, which affect the strength of intermolecular forces. For example, the addition of heat to a substance causes its particle's thermal energy (or the energy of motion) to increase, overcoming the attractive intermolecular forces between them. A solid melts when its temperature rises to the point at which the particles vibrate fast enough to move out of their fixed positions. This phase transition is called melting, and the point at which it occurs is the solid's melting point. As the temperature increases further, the particles move faster until they finally escape into the gaseous state. This is vaporization, and the point at which it occurs is the liquid's boiling point.

The phase transition point and the energy change associated with the transition depend on the intermolecular forces that exist in the substance. At a given pressure, substances with stronger intermolecular forces require more energy to overcome them and, therefore, undergo phase changes at higher temperatures. The energy required to cause the complete phase transition of one mole of a substance without a change in temperature is called the molar heat or molar enthalpy of that transition. For example, the energy required to vaporize one mole of a liquid is called the molar enthalpy of vaporization.

Transitions that occur by absorbing energy are exothermic, and their enthalpy values are negative. On the other hand, transitions that occur by releasing energy are endothermic, and their enthalpy values are positive. For example, while the molar enthalpy of vaporization is positive, the molar enthalpy of condensation is negative.

Because a substance transforms from one phase to another molecule by molecule, during a phase transition, the two phases coexist; and the temperature of the substance stays constant, despite the continuous supply of heat. After the transition of the bulk completes, the temperature of the substance rises.

When phase transitions occur in a closed system, the opposing transitions occur at equal rates, leading to a state of dynamic equilibrium.