Colligative Properties change due to the presence of a solute.  The magnitude of the change is due to the number of particles in the solution and not their chemical identity.  Examples of properties that fall under this category are the vapor pressure, melting and boiling points, and osmotic pressure.  Changes of state depend on vapor pressure, so the presence of a solute affects the freezing point and boiling point of a solvent. In a solution only a fraction of the molecules with energy in excess of the intermolecular force are nonvolatile solute molecules. This reduction in the number of solvent molecules with energies above the intermolecular force lowers vapor pressure of the solution than the pure solvent at all temperatures. The presence of solute particles interferes with the crystallization process and thus the normal melting point is lower; this means more energy needs to be removed to favor crystallization or the formation of a solid. For example, the normal boiling point of a liquid occurs at the temperature where the vapor pressure is equal to 1 atmosphere. A nonvolatile solute elevates the boiling point of the solvent.  The magnitude of the boiling-point elevation depends on the concentration of the particles of solute.  When a solute is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solvent. The equation for freezing-point depression is similar to that for boiling-point elevation.

Osmotic Pressure occurs when a solution and a pure solvent are separated by a semipermeable membrane, which does not allow the solute molecules to pass through only allowing the solvent to cross the membrane. The flow of solvent (in living organisms typically water) is called osmosis. The excess pressure compared to pure solvent is called osmotic pressure. Solutions that have ideal osmotic pressure (at equilibrium) are said to be isotonic. When a solution in contact with pure solvent across a semipermeable membrane is subjected to external pressure greater than the osmotic pressure, then reverse osmosis occurs.

Osmolarity or osmotic concentration is the number of osmoles per Liter of solution (osmol/L or Osm/L). Molarity measure the moles of solute per unit volume of solution, osmolarity measure the number of particles or number of osmoles per unit of volume. When compounds dissociate, the number of osmoles increases. Ionic compounds such as sodium chloride (NaCl) can dissociate in solution into their positively and negatively charged ions. When NaCl dissociates into Na⁺ and Cl⁻ ions, for every 1 mole of NaCl in solution, there are 2 osmoles of solute particles (i.e., a 1 mol/L NaCl solution is a 2 osmol/L NaCl solution). Both sodium and chloride ions affect the osmotic pressure of the solution. Another example is magnesium chloride (MgCl₂), which dissociates into Mg²⁺ and 2Cl^– ions. For every 1 mole of MgCl₂ in the solution, there are 3 osmoles of solute particles. Nonionic compounds do not dissociate, and form only 1 osmole of solute per 1 mole of solute. For example, a 1 mol/L solution of glucose is 1 osmol/L. Multiple compounds in solution will contribute to the osmolarity of a solution. For example, a 4 Osm solution might consist of: 4 moles glucose, or 2.0 moles NaCl, or 2 mole glucose + 1 mole NaCl, or any other type of combination.