The internal energy of a substance—the total kinetic energy of all its molecules and the potential energy of their associated forces—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 the gaseous state, the lowest in the solid state, and intermediate in the liquid state. Phase transitions are caused by changes in physical conditions, such as temperature and pressure, that affect intermolecular interactions. For example, the addition of heat to a substance causes the thermal energy of its particles to increase. When the total heat added causes the particles in a solid to vibrate fast enough to overcome their intermolecular attraction, they 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. During a phase transition, molecules exist in both phases simultaneously and the temperature of the substance stays constant despite the continuous influx of heat. After the transition of the bulk of the substance is completed, the temperature of the substance rises.
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 undergo phase transitions, so these transitions occur at higher temperatures than substances with weaker intermolecular interactions. 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. Reactions or transitions that absorb energy are endothermic, and reactions or transitions that release energy are exothermic. Reactions that absorb energy have positive enthalpy values while reactions that release energy have negative enthalpy values. For example, while the molar enthalpy of vaporization is positive, the molar enthalpy of condensation is negative, so vaporization is an endothermic process while condensation is exothermic.
