13.3 Phase Transitions
Learning Objectives
- Define phase transitions and phase transition temperatures
- Explain the relation between phase transition temperatures, intermolecular attractive forces, and kinetic energy
- Describe the processes represented by typical heating and cooling curves, and compute heat flows and enthalpy changes accompanying these processes
We witness and utilize changes of physical state, or phase transitions, in a great number of ways. As one example of global significance, consider the evaporation, condensation, freezing, and melting of water. These changes of state are essential aspects of our earth’s water cycle as well as many other natural phenomena and technological processes of central importance to our lives. In this module, the essential aspects of phase transitions are explored.
Vaporization and Condensation
When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally collide with the surface of the condensed phase, and in some cases, these collisions will result in the molecules re-entering the condensed phase. Figure 13.3a shows a simple phase change diagram. The change from the gas phase to the liquid is called condensation. When the rate of condensation becomes equal to the rate of vaporization, neither the amount of the liquid nor the amount of the vapour in the container changes. The vapour in the container is then said to be in equilibrium with the liquid. Keep in mind that this is not a static situation, as molecules are continually exchanged between the condensed and gaseous phases. Such is an example of a dynamic equilibrium, the status of a system in which reciprocal processes (for example, vaporization and condensation) occur at equal rates. The pressure exerted by the vapour in equilibrium with a liquid in a closed container at a given temperature is called the liquid’s vapour pressure (or equilibrium vapour pressure). The area of the surface of the liquid in contact with a vapour and the size of the vessel have no effect on the vapour pressure, although they do affect the time required for the equilibrium to be reached. We can measure the vapour pressure of a liquid by placing a sample in a closed container, like that illustrated in Figure 13.3b, and using a manometer to measure the increase in pressure that is due to the vapour in equilibrium with the condensed phase.
The chemical identities of the molecules in a liquid determine the types (and strengths) of intermolecular attractions possible; consequently, different substances will exhibit different equilibrium vapour pressures. Relatively strong intermolecular attractive forces will serve to impede vaporization as well as favouring “recapture” of gas-phase molecules when they collide with the liquid surface, resulting in a relatively low vapour pressure. Weak intermolecular attractions present less of a barrier to vaporization, and a reduced likelihood of