Assume a scenario where two immiscible liquids are in a closed flask, with one floating on top of the other. In this situation, if the system is kept still, only the vapor pressure of the liquid on top will be measured because it's the only one in contact with the vapor phase.
However, if the mixture is agitated or stirred, both liquids will be broken up into drops, and at any given time, there will be drops of both liquids on the surface. This means that both liquids contribute to the overall vapor pressure of the mixture.
When the mixture is agitated, both liquids are in equilibrium with their vapors, and the total vapor pressure is simply the sum of the individual vapor pressures. This is described by the equation:
Total vapor pressure = p°A + p°B
Where p°A and p°B refer to the saturated vapor pressures of the pure liquids A and B, respectively. This equation is independent of the amount of each liquid present, as long as there's enough of each for both to exist in equilibrium with their vapor.
While the concept of vapor pressure applies similarly to both immiscible and miscible liquids, there are important distinctions between the behavior of these two types of mixtures when it comes to vapor pressure and boiling points.
In both cases, the vapor pressure of a mixture is determined by the partial pressures of each component. However, in immiscible liquids, the components do not mix at a molecular level, so each component maintains its own distinct vapor pressure regardless of the presence of the other. In contrast, miscible liquids form a homogeneous mixture, meaning the individual components dissolve into each other at a molecular level, resulting in a combined vapor pressure that is not simply the sum of the individual vapor pressures. Instead, it's determined by the vapor pressure of the mixture as a whole, which can be higher or lower than the sum of the individual vapor pressures depending on the interactions between the components.
