The more solute particles there are in solution, the lower the mole fraction of the solvent would be,
and hence the lower the vapor pressure. One limiting case is the trivial scenario where only the
solute is present; without any solvent, XA is zero, and so the vapor pressure of A is zero by the
equation, which certainly makes sense because there simply isn’t any solvent around to exert a
vapor pressure. In the other extreme, when no solute is present, the system is composed entirely of
solvent A; its mole fraction is therefore one and its vapor pressure would be the same as that of pure
A. In between these two cases, Raoult’s law states that the vapor pressure is linearly proportional to
the mole fraction of the solvent.
It should be pointed out that even though we have introduced Raoult’s law in the study of colligative
properties, it is not limited to the solvent-nonvolatile solute systems on which we have focused. In a
solution with several volatile components (a mixture of benzene and toluene, for example), Raoult’s
law states that the vapor pressure of each component is proportional to its mole fraction in the
solution below.
The sum of all the mole fractions has to equal one. The total vapor pressure over the solution, then,
is the sum of the partial vapor pressures of each component:
Ptot = PA + PB + PC + ... = XAP°A + XBP°B + X (^) C P°C + ...
This last result is simply an application of Dalton’s law of partial pressures.
Raoult’s law is actually only an idealized description of the behavior of solutions, and holds only
when the attraction between molecules of the different components of the mixture is equal to the
attraction between the molecules of any one component in its pure state. When this condition does
not hold, the relationship between mole fraction and vapor pressure will deviate from Raoult’s law.
Solutions that obey Raoult’s law are called ideal solutions, much in the same way that gases obeying
PV = nRT are called ideal gases.