"Based on Figure 13.18, you might think that the reason volatile solvent molecules in a solution are less likely to escape to the gas phase, compared to the pure solvent, is that the solute molecules are physically blocking the solvent molecules from leaving at the surface. This is a common misconception. Design an experiment to test the hypothesis that solute blocking of solvent vaporization is not the reason that solutions have lower vapor pressures than pure solvents."
The explanation (in the main chapter, not as an answer to this question) given in Chemistry: The Central Science (13th Edition) by Brown, LeMay, and Bursten (abbreviated as CTCS) is: "A solution consisting of a volatile liquid solvent and a nonvolatile solute forms spontaneously because of the increase in entropy that accompanies their mixing. In effect, the solvent molecules are stabilized in their liquid state by this process and thus have a lower tendency to escape into the vapor state. Therefore, when a nonvolatile solute is present, the vapor pressure of the solvent is lower than the vapor pressure of the pure solvent, as illustrated in ▲ Figure 13.18."
However, this is, I believe, the only textbook which gives this explanation for lowering in vapour pressure on adding a non-volatile solute. The other (reputed) textbooks I've read, like Zumdahl and NCERT Grade XII Chemistry, all give the explanation that CTCS calls a "common misconception". So, which explanation is correct? If possible, please give an experiment also.