Question

A solution contains 0.0442 mol of ethanol and 0.715 mol of water at 25 oC. The vapor pressures of pure ethanol and pure water at 25 oC are 44.6 torr and 23.8 Torr, respectively.

a) Assuming ideal behavior, calculate the total vapor pressure above the solution.

b) If the actual vapor pressure above the solution is found to be 28.3 Torr, is the solution behaving ideally? If not, what interactions are favored (the solute-solute and solvent-solvent interactions, or the solute interactions)?

Answer 1

(a)

Vapor pressure of solution = (vapor pressure of pure ethanol x mole fraction ethanol) + (vapor pressure of water x mole fraction cyclohexanone).


mole fraction ethanol = 0.0442 / (0.0442 + 0.715) = 0.0442 / 0.7592 = 0.0582
mole fraction water = 0.715 / (0.0442 + 0.715) = 0.9418

vapor pressure solution = 44.6 x 0.0582 + 23.8 x 0.9418 = 2.596 + 22.415 = 25.011 torr

(b)

When the solution behave ideally the vapour pressure is 25.011 torr.

All the three interactions are favored. Since

In order for a solute to be soluble in a particular solvent, three things need to be taken in consideration.

(1) The intermolecular forces holding the solvent molecules together must be broken to make room for the solute. (bond breaking always takes energy).

(2) The intermolecular forces holding the solute together must be broken. (Again, this requires energy).

(3) The solute and solvent can interact through whatever forces are available. (bond formation gives off energy), and is called solvation energy. If the energy needed in the first two steps is not too great and the solvation energy is sufficiently large, the overall energy will be negative or slightly positive, a solution will be formed.