In a solution, there are three major attractive intermolecular forces: attractions between the solvent molecules, attractions between the solute molecules, and attractions between the solute and the solvent molecules.

If the strength of each of the three types of interactions are similar in magnitude, the solution is called an ideal solution.

An ideal solution obeys Raoult’s law at all concentrations. 

For an ideal solution with two volatile components, such as toluene and benzene, the partial vapor pressure of each component will be given by Raoult’s law as a product of the vapor pressure of the pure component and its mole fraction.

In a given solution, the mole fraction for toluene is 0.4 and the mole fraction for benzene is 0.6. As the vapor pressures of the pure toluene and pure benzene are 22 and 75 torr, respectively, the partial pressures of toluene and benzene in this solution will be 8.8 and 45 torr, respectively.

The total vapor pressure is the sum of the partial pressure of each component and is equal to 54 torr.

For such an ideal solution, a plot of vapor pressure against mole fraction yields a straight line.

When the intermolecular forces within a solution are not uniform, the solution deviates from Raoult’s law and is called non-ideal.

If the solvent-solute interactions in a solution are weaker than the solvent–solvent interactions, as in the case of a benzene and methanol solution, the solute will allow more solvent particles to escape into the gaseous state than in the pure solvent.

Thus, vapor pressure would tend to be greater than that predicted by Raoult’s law. Such solutions show a positive deviation from Raoult’s law.

Conversely, in a solution with strong solute-solvent interactions, the solute will prevent the solvent from vaporizing and the vapor pressure of the solution will be less than that predicted by Raoult’s law. 

This is observed in an aqueous solution of acetone and chloroform, where strong hydrogen bonding between the two leads to a negative deviation from Raoult’s law.