I was studying colligative properties of dilute solutions, and one of the properties was freezing point depression. The confusion I had was that, if vapour pressure decreases, it should be easier to reach the freezing point (point where vapour pressure of solid and liquid are equal). So I thought the freezing point will increase, but freezing point actually decreases. Is there any logical explanation?
Consider graphs of the vapor pressure of the liquid and of the solid, both as a function of temperature. Intuitively, both graphs increase with temperature, and intersect at the freezing point. By lowering the vapor pressure of the liquid, we lower the corresponding vapor pressure curve and hence the temperature at the point of intersection.
I think it of this way. There is a lot of intermolecular space between individual molecules of a solvent. When we decrease the temperature, the individual molecules come closer, lose their heat of fusion and eventually the intermolecular spaces goes on decreasing until they became a tightly packed solid.
Adding a solute hinders the process as the solute particles, which are different in size and geometry from the solvent particles, occupies these interstitial spaces and prevents solvent molecules from coming closer. Hence we need a much lower temperature than before to freeze my solution as at lower temperatures, the solvent-solute particle interaction gets relatively weakened.