Final answer:
Colligative properties of solutions, such as boiling point elevation, freezing point depression, and osmotic pressure, depend solely on the number of solute particles dissolved, irrespective of their nature. These properties result from the disruption of the solvent's natural physical properties by the solute particles.
Step-by-step explanation:
Colligative properties of solutions depend ONLY on the number of solute particles present, not their chemical identity. This means that factors like boiling point elevation, freezing point depression, and osmotic pressure changes in a solution are influenced by how many particles of a solute are dissolved rather than what those particles are. For example, a 0.001 mol/kg aqueous solution of any non-electrolyte should result in the same adjustment to the solution’s boiling point, freezing point, and osmotic pressure.
These properties arise because the presence of solute particles disrupts the solvent's natural properties. In freezing point depression and boiling point elevation, the addition of solute lowers the solvent's vapor pressure, altering the temperatures at which the solution will transition between states (solid/liquid for freezing point, liquid/gas for boiling point). For electrolyte solutions that dissociate into ions, the effect on colligative properties will be more pronounced compared to non-electrolytes because more particles are created per molecule of solute.
The relationship between vapor pressure and mole fraction is described by Raoult's law. In pure solvents, vapor pressure is entirely due to the solvent molecules; however, with added solute, the vapor pressure over the solution is lower because the solute occupies some of the surface area that would otherwise be occupied by the solvent. The degree to which a solute affects the colligative properties of a solution can also be described by the van't Hoff factor, which takes into account the dissociation of solutes in a solution.