Colligative properties
Encyclopedia : C : CO : COL : Colligative properties
In chemistry, colligative properties are factors that determine how the properties of a bulk liquid solution change depending on the concentration of the solute in the bulk solution. It is important to note that this is not relative concentrations, as one sees in azeotropes for example, but bulk concentrations where the solvent volume is held constant.
The identity of the solute(s) has no first order importance in the consideration of colligative properties, and the degree to which their manifestation can be observed and measured is affected only by the number of solute particles in the solution. However, the identity of solute may affect the number of effective molecules in solution; for example, hydrogen bonding. These types of considerations are approximate however, since in general, colligative models make assumptions of ideality wherein molecular interactions are neglected.
The four colligative properties are:
- Vapor pressure: The change in vapor pressure where the solute is less volatile than the solvent is regulated by Raoult's law, which states that the pressure is equal to the mole fraction of the solvent times the vapor pressure of pure solvent: P=Xsolvent*P°. This holds truest for ideal solutions.
- Freezing-point depression: The presence of a solute decreases the freezing point as compared to a pure solvent. The exact change (ΔT) can be calculated as van 't Hoff factor (i) of the solute multiplied by its molality (m) multiplied by the freezing point depression constant of the solvent (Kf): ΔT=i(-Kf)m. Alternatively, it can be calculated as the total molality of all solutes in solution times the depression constant: ΔT=Kf∑m.
- Boiling-point elevation: Because of the lowered vapor pressure, the boiling point of a solution is elevated as compared to the pure solvent. The change in boiling point (ΔT) can be calculated in the same way as the change in freezing point, except a different constant is used: the van 't Hoff factor (i) of the solute multiplied by its molality multiplied by the boiling point elevation constant of the solvent (Kb): ΔT=iKbm. ΔT can also be calculated using the sum of all molalities: ΔT=Kf∑m.
- Osmotic pressure: The presence of solute can cause pressure to be exerted across a permeable membrane according to an equation quite similar to the ideal gas law: π is the osmotic pressure, n is the number of moles of solute, R is the ideal gas constant, T is the absolute temperature in kelvins, and V is the volume: π[=(nRTi)/V]
In (1), the two columns of pure solvent (blue) under the same pressure. When solute (green) is added to the right column (2), osmotic pressure is exerted and solvent flows through the permeable membrane (red) to the right side.
From Wikipedia, the Free Encyclopedia. Original article here. Support Wikipedia by contributing or donating.
All text is available under the terms of the GNU Free Documentation License See Wikipedia Copyrights for details.