- Online Text
- 1. Introduction
- 2. What Is a Solution?
- 3. Solutions and Solubility
- 4. Solution Concentrations
- 5. Analyzing Solutions
- 6. Raoult's Law
- 7. Henry's Law
- 8. Colligative Properties—Vapor Pressure and Osmosis
- 9. Colligative Properties—Freezing and Boiling
- 10. Separation and Purification
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 8: Colligative Properties—Vapor Pressure and Osmosis
Chemists typically work with solutions that are dilute, meaning that they consist mostly of solvent with only a little solute. In these cases, the mole fraction of the solvent is high and may be close to one. Of course, it is not actually one, because then we would have a pure substance instead of a solution.
Many physical properties of solutions depend only on the concentration or amount of the solute in the solution, not on the identity of the solute. In these, the concentration can be measured as a mole fraction, molarity or molality. These properties are called "colligative properties" and include:
- Vapor pressure depression
- Osmotic pressure
- Boiling point elevation
- Freezing point depression
These four properties are bound together by the fact that they each vary in proportion to the amount of solute and not the type of solute present. They are also closely related to chemistry behind Raoult's Law. Let's look separately at the first two in this section.
Vapor Pressure Depression
Figure 8-12. Vapor Pressure Depression
The vapor pressure above a pure volatile liquid will always be higher than a solution of that liquid with some solute dissolved in it.
© Science Media Group.
From Raoult's Law, we know that the partial pressure of a gas above a mixture is proportional to the mole fraction of that component in the liquid. If the solution consists of a solute that is nonvolatile (that is, it doesn't evaporate) and a solvent that is volatile (it will produce vapor), then the only vapor above the liquid will be that of the solvent. An example would be an aqueous solution of sodium chloride: The sodium chloride will not evaporate to form a gas above the solution, but the water will.
Because the solvent has a solute dissolved in it, the mole fraction of the solvent is less than one, so the partial pressure is less than the vapor pressure of pure solvent. The amount by which it is less is called the "vapor pressure depression." The more concentrated the solute, the greater the vapor pressure depression. (Figure 8-12) This means that if we take a glass of water from the tap and some salt water from the ocean and leave them sitting out, the salt water will evaporate much more slowly than the pure water. As the salt water evaporates, the concentration of the salt will rise, the vapor pressure of the water will keep falling, and it will take longer and longer to evaporate the same amount of water.
When a membrane or similar barrier allows the passage of some molecules but not others, it is called a "semipermeable membrane." If a semipermeable membrane is placed in a solution, generally the solvent flows freely across the membrane and the solute is blocked. When this happens, solvent will flow through the membrane toward the side on which there is a higher concentration of solute. The sidebar on how cells behave in solutions shows some of the interesting effects of osmotic pressure. (See the Osmotic Pressure and Cells sidebar.)
In Figure 8-13, we have placed solvent in a jar and set a glass tube in it with a semipermeable membrane at the bottom. Some solute particles are dissolved in the solvent inside the tube, where they are confined because they cannot pass through the semipermeable membrane. Because the solution in the tube is more concentrated, solvent will flow through the membrane into the tube to try to dilute it. This won't happen forever, just until the liquid reaches a certain height. The height of this column of water can be related to what we call its osmotic pressure, which has forced the liquid up to that height.
Figure 8-13. Osmotic Pressure
A. When solvent is free to flow across a semipermeable membrane, it will flow into the solution to dilute it somewhat. B. We can counteract this flow by placing a plunger on the liquid in the tube and pressing down. C. Pushing down on the plunger will force pure water out through the membrane.
© Science Media Group.
This pressure can affect cells in the body making them swell or shrink, but it has a practical industrial use: reverse osmosis. If a pressure higher than the osmotic pressure of solution is applied to it, as by pushing down on a plunger, pure water can be forced out of the semipermeable membrane. This is one of the common ways, not involving filters or evaporation, to purify water.