in a non ideal solution with +ve deviation Pa+Pb=Ps Ps[non ideal sol]>Ps[ideal sol.] so, why Pa[non ideal sol]>Pa[ideal sol.] and Pb[non ideal sol]>Pb[ideal sol.] ? pls ans acc. to graphs also i m so confused ?

To understand why in a non-ideal solution with positive deviation, the partial pressures of the components are greater than those in an ideal solution, we need to delve into the concepts of vapor pressure, Raoult's Law, and how intermolecular forces play a role in these deviations.

Understanding Vapor Pressure and Raoult's Law

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid or solid phase. In an ideal solution, the vapor pressure of each component is directly proportional to its mole fraction in the solution. This relationship is described by Raoult's Law, which states:

• For Component A: P_A = X_A * P_A⁰
• For Component B: P_B = X_B * P_B⁰

Here, P_A and P_B are the partial pressures of components A and B, X_A and X_B are their mole fractions, and P_A⁰ and P_B⁰ are their vapor pressures in pure form.

Positive Deviation from Ideal Behavior

In a non-ideal solution with positive deviation, the interactions between the molecules of different components are weaker than those between the molecules of the same component. This weaker interaction leads to an increase in the tendency of the molecules to escape into the vapor phase, resulting in higher vapor pressures than predicted by Raoult's Law.

Why Are Partial Pressures Higher?

Now, let’s break down why P_A[non-ideal] > P_A[ideal] and P_B[non-ideal] > P_B[ideal].

• Weaker Intermolecular Forces: In a non-ideal solution, the molecules of different components do not interact as strongly as they do in their pure states. This means that when they are mixed, they experience less cohesive force, allowing more molecules to escape into the vapor phase.
• Higher Mole Fraction Contribution: Since the vapor pressures are higher, the mole fractions of the components in the vapor phase will also be higher. This results in a greater partial pressure for each component compared to what would be expected in an ideal solution.

Graphical Representation

To visualize this, consider a graph where the x-axis represents the composition of the solution (mole fraction of component A) and the y-axis represents the vapor pressure. In an ideal solution, the line representing the total vapor pressure would be a straight line, reflecting the linear relationship dictated by Raoult's Law. However, for a non-ideal solution with positive deviation, the curve would lie above this line, indicating that the actual vapor pressures are higher than those predicted by Raoult's Law.

Example to Illustrate

Imagine a solution of ethanol and water. Ethanol has weaker hydrogen bonds compared to water. When mixed, the ethanol molecules disrupt the hydrogen bonding network of water, leading to a situation where both components can escape into the vapor phase more readily than they would in their pure states. Thus, the partial pressures of both ethanol and water in the vapor phase are higher than what would be expected if they behaved ideally.

In summary, the positive deviation in non-ideal solutions arises from weaker intermolecular forces and results in higher partial pressures for each component compared to their ideal counterparts. This understanding is crucial for predicting the behavior of mixtures in various chemical processes.