As a metal such as lead melts, what happens to (a) the average kinetic energy of the atoms and (b) the average distance between the atoms?

Step 1 of 5) The vapor–pressure curves for the liquid and solid phases meet at the triple point. (Section 11.6) In Figure 13.21 we see that the triple-point temperature of the solution is lower than the triple-point temperature of pure liquid because the solution has a lower vapor pressure than the pure liquid. Phase diagram illustrating freezing-point depression. The black lines show the pure solvent’s phase equilibria curves, and the blue lines show the solution’s phase equilibria curves. The freezing point of a solution is the temperature at which the first crystals of pure solvent form in equilibrium with the solution. Recall from Section 11.6 that the line representing the solid–liquid equilibrium rises nearly vertically from the triple point. It is easy to see in Figure 13.21 that the triple-point temperature of the solution is lower than that of the pure liquid, but it is also true for all points along the solid–liquid equilibrium curve: the freezing point of the solution is lower than that of the pure liquid. Like the boiling-point elevation, the change in freezing point ∆Tf is directly proportional to solute molality, taking into account the van’t Hoff factor i: