• Consider the molecules of a liquid A and its vapour are in random motion. If the liquid is in an open container, the molecules in the vapour phase spread. Consequently, more and more molecules enter the vapour phase from the liquid phase till the whole liquid evaporates.
• When the evaporation occurs in a closed space, molecules enter the vapour phase from the liquid phase and vice versa and a dynamic equilibrium is established between the two phases at the temperature concerned. At this equilibrium, the rate of evaporation is equal to the rate of condensation.
A(l) ⇌ A(g)
However, the microscopic changes are not observable. If the temperature remains unchanged the pressure exerted by the vapour at equilibrium remains constant and it is the saturated vapour pressure at the particular temperature. This itself is an equilibrium constant.
• The kinetic energy of molecules depends on the temperature. When the temperature rises, the kinetic energy increases and the molecules with high kinetic energy overcome the intermolecular forces and move into the vapour phase. The higher the intermolecular attractions, the lesser the tendency of the molecules to pass into the vapour phase.
• As more and more molecules move into the vapour phase, the vapour pressure increases. Hence with increasing temperature, the vapour pressure also increases. When the temperature is decreased, the kinetic energy of the molecules in the vapour phase decreases with the result that more molecules return to the liquid phase and the vapour pressure drops.
• When the liquid is heated, the temperature rises and at a certain temperature the saturated vapour pressure becomes equal to the (external) atmospheric pressure. At this temperature, the liquid boils and therefore, it is the normal boiling point of the liquid.
• The increase in vapour pressure of liquids with temperature is not linear . Moreover, different liquids have different intermolecular attractive forces and hence different volatilities. Therefore, the temperatures at which their s.v.p. becomes equal to the atmospheric pressure are different, so they boil at different temperatures
• Consider the liquid-vapour equilibrium established in a closed system. With rise in temperature the equilibrium is disturbed and more molecules pass into the vapour phase from the liquid phase establishing new equilibria increasing the amount in the gas phase. Finally a stage is reached where only the vapour exists.
• This vapour can be liquefied by compression. However, the tendency of the vapour to liquefy decreases with increasing temperature. Hence, a gas has a characteristic minimum temperature, called the critical temperature, above which the gas cannot be liquefied by increasing pressure, no matter how high the pressure would be.
The minimum pressure required to bring about liquefaction at the critical temperature is known as the critical pressure and the volume of one mole of the substance at the critical temperature and critical pressure is known as the critical volume.
• The following phase diagram of water shows the variation of vapour pressure of water with temperature (curve TC) and vapour pressure of ice with temperature (curve AT). Line BT shows the temperatures and pressures at which ice and liquid water are in equilibrium.
At the point T ice, water and vapour exist together in equilibrium, and this point is known as the triple point. The curve TC ends at the critical temperature (647.4 K) of water. Above this temperature only the vapour phase can exist.
