Some non-ideal solutions show positive deviation from Raoult's law. Consider a solution of two components A and B. If A-B interactions in the solution are weaker than the A-A and B-B interactions in the two liquids forming the solution, then the escaping tendency of molecules A and B from the solution become more than in pure liquids. The total vapour pressure will be greater than the corresponding vapour pressure as expected on the basis of Raoult's law. This type of behaviour of solution is called positive deviation from Raoult's law. Tbe boiling point of such solutions are lowered. Mathematically,

P_A < p^o_AX x_A
PB < p^o_B x x_B
The total vapour pressure is less than PA + PB
PA + P_B
P < P^o_AX x_A + P⁰_B x_B
P₁ = P_A
P₂ = P_B
Examples of solutions showing positive deviations:
(i) Etbyl alcobol and water
(ii) Benzene and acetone
(iii) Ethyl alcohol and cyclobexane
(iv) Carbon tetrachloride and chlorofor

Ideal solutions: The solutions which obey Raoult's law over the entire range of concentration are known as ideal solutions. Ideal solutions are formed by mixing the two components which are identical in molecular size, in structure and have almost identical intermolecular forces.
Examples:
(i) Benzene and toluene
(ii) n-Hexane and n-Heptane
(iii) Chlorobenzene and bromobenzene.
Characteristics:
(i) They must obey Raoult's law.
(ii) ΔH mixing should be zero.
(iii)ΔV mixing should be zero, i.e. volume change on mixing is zero.

Non-ideal solutions : The solutions which do not obey Raoult's law are called non-ideal solutions. In case of non-ideal solutions there is change in volume and heat energy when the two componnts are mixed.
Characterstics :
(i) They does not Rault's law.
(ii)△V = mix ≠ 0
(iii) △H mix ≠ 0
Behaviour of Ideal Solutions: A plot of P_I or P₂ versus the mole fraction x₁ and x2 for an ideal solution gives a linear plot. These Lines (I and II) pass through the points and respectively when x _I and x₂ is equal to unity. Similarly the plot (Line III) of P_total versus x₂ is also linear. The minimum value of P_total is P_total and the maximum value is P₂⁰, assuming that component I is less volatile than component 2, i.e. P1 O < P2 O

According to van't Hoff equation
\(\pi =CRT\)
\(C=\cfrac { n }{ V } \)
Here n = number of moles of solute dissolved in 'V' litre of the solution
\(\pi =\cfrac { n }{ V } .RT\Rightarrow \pi V=nRT\)
If the solution is prepared by dissolving WB of the non-volatile solute in WA g of solvent, then the number of moles of 'n' is
\(n=\cfrac { { W }_{ B } }{ M_{ B } } \) where, M_B = molar mass of the solute
Substituting n value, we get
\(\pi =\cfrac { { W }_{ B } }{ V } .\cfrac { RT }{ { M }_{ B } } \)
\(\therefore\)\({ M }_{ B }=\cfrac { { W }_{ B } }{ V } .\cfrac { RT }{ \pi } \)

(i) The measurement of relative lowering of vapour pressure can be used to determine the molar mass of a non-volatile solute.
(ii) A known mass of the solute is dissolved in a known quantity of solvent. The relative lowering of vapour pressure is measured experimentally.
(iii) According to Raoult's law, the relative lowering of vapour pressure is
\(\cfrac { { P }_{ solvent }^{ o }-{ P }_{ solution } }{ { P }_{ solvent }^{ o } } =xB\)
W_A = weight of solvent, WB = weight of solute
M_A= Molar mass of solvent, MB = molar mass of solute
\(\therefore { x }_{ B }=\cfrac { { n }_{ B } }{ { n }_{ A }+{ n }_{ B } } \)
where n_A= number of moles of solvent, n_B = number of moles of solute.
For dilute solution, n_A>> nB, n_A + n_B≈nA.
Then ,\({ x }_{ B }=\cfrac { { n }_{ B } }{ { n }_{ A } } \)
Number of moles of solvent and solute are
\({ n }_{ A }=\cfrac { { W }_{ A } }{ { M }_{ A } } .{ n }_{ B }=\cfrac { { W }_{ B } }{ { M }_{ B } } \)
\(\therefore { x }_{ B }=\cfrac { \frac { { W }_{ b } }{ { M }_{ B } } }{ \frac { { W }_{ A } }{ { M }_{ A } } } \)
Thus, relative lowering of vapour pressure = \(\cfrac { \frac { { W }_{ B } }{ { M }_{ B } } }{ \frac { { W }_{ A } }{ { M }_{ A } } } \)
Relative lowering of vapour pressure = \(\cfrac { { P }^{ o }-P }{ { P }^{ o } } \) 
\(\cfrac { { P }^{ o }-P }{ { P }^{ o } } =\cfrac { { W }_{ B }-{ M }_{ A } }{ { W }_{ A }\times { M }_{ B } } \)
From the above equation, molar mass of the solute MB can be calculated using the known values of W_A'W_B' M_Aand the measured relative lowering of vapour pressure.

(i) Solute-solvent interactions: For an ideal solution, the interaction between the solvent molecules (A-A), the solute molecules (B-B) and between the solvent and solute molecules (A-B) are expected to be similar. If these interactions are dissimilar, there will be a deviation from ideal behaviour.
(ii) Dissolution of solute: When a solute present in a solution dissociates to give its constituent ions, the resultant ions interact strongly with the solvent and causes deviation from Raoult's law. e.g., KCI in water deviates from ideal behaviour due to dissociation as K+ and Cl" ion which form strong ion-dipole interaction with water molecules.
(iii) Association of solute: Association of solute molecules can also cause deviation from ideal behaviour. For example in solution acetic acid exists as a dimer by forming intermolecular hydrogen bonds and hence deviates from Raoult's law.
(iv) Temperature: An increase in temperature of the solution increases the average kinetic energy of the molecules present in the solution which cause decrease in the attractive force between them. As result, the solution deviates from Raoult's law.
(v) Pressure: At high pressure, the molecules tends to stay close to each other and therefore there will be an increase in their intermolecular attraction. Thus a solution deviates from Raoult's law at high pressure.
(vi) Concentration: When the concentration is increased by adding solute, the solvent-solute interaction becomes significant. This causes deviation from Raoult's law.