Phase Transitions

Substances typically exist in one of three phases: solid, liquid, or gas. Transitioning from one phase to another significantly changes how ordered and tightly held the molecules are.

Molecules transition between different phases when their internal energy allows them to be in either state. This depends on factors like the strength of the intermolecular forces in the more-condensed state and the pressure exerted on the substance.

Temperature is a reflection of internal energy, so phase transition points are often described in terms of temperature at a certain pressure.

For instance, compare water and acetone. While water exhibits strong hydrogen bonds, acetone molecules display weak dipole-dipole forces.

Due to the stronger forces of attraction, more heat is required to turn water into steam. This explains why at any given pressure, the boiling point of acetone is lower than that of water.

Phase transitions happen molecule by molecule, so the phases coexist during the transition.

Until the transition of the bulk is complete, the temperature does not change even though heat is flowing to or from the substance.

A similar observation is made when heat is supplied to water. The temperature of water rises until it reaches its boiling point, at which the two phases - liquid and gas - coexist. Additional heating does not increase the temperature of the liquid water beyond its boiling point; instead, it only causes more rapid boiling.

The change in energy required for one mole of a substance to completely undergo that transition without a change in temperature is called the molar heat or molar enthalpy of that transition.

If a substance absorbs heat to undergo a transition, the enthalpy of the transition is positive, making it an endothermic process. Transitions in which the substance loses heat have negative enthalpy values, making them exothermic.

If a substance is held at a transition point in a closed system, the opposing transition processes will reach a state of dynamic equilibrium.