The colligative properties of a solution are those properties which depend upon the number of molecules in it, rather than on what sort of molecules they are. It’s all determined by the ratio of number of solute particles to the number of solvent particles. There could be any particles whatsoever.
The colligative properties are:
- Vapour pressure
- Boiling point
- Freezing point
- Osmotic pressure
Vapour pressure, boiling point and freezing point (as well as of their respiratory-related matters such as humidity) are discussed in a separate chapter on the physics of humidity and evaporation.
Introducing solutes into your solvent changes the colligative properties. Doesn’t matter what solute- these properties don’t discriminate. The number of particles is the most important issue.
- Boiling point increases in proportion to the molar concentration of the solutes
- Freezing point decreases in proportion to the molar concentration of the solutes
- Vapour pressure is dependent on the vapour pressure of each chemical component and the mole fraction of that component in the solution (Raoult’s law)
Equilibrium Vapour Pressure
This is a measure of the evaporation rate of a liquid.
Picture a closed canister with some liquid.
That liquid at all times has the tendency to evaporate.
Its vapour also has the constant tendency to condense.
Left alone, these two will reach an equilibrium, where the rate of evaporation matches the rate of condensation.
The pressure in the vapour at this equilibrium is the equilibrium vapour pressure.
The vapour pressure increases as temperature increases (but the relationship is non-linear).
Like all pressures, it can be measured in mmHg. At 760mmHg atmospheric pressure and 25 C° temperature the vapour pressure of water is about 25mmHg; at body temeprature it is about 47 C°.
In a solution with multiple solutes, the vapour pressure will be depressed.
Each solute’s individual vapour pressure contributes to the total vapour pressure, in proportion to its mole fraction. The mole fraction is the proportion of the total molarity of solutes this solute contributes. So the more solutes you add, the lower the individual solute mole fraction and the lower the contributed vapour pressure.
That’s confusing. The bottom line is, the more solutes you add to a solution, the lower the vapour pressure.
Atmospheric Pressure Boiling Point
This is the temperature at which vapour pressure equals atmospheric pressure. Above this temperature, bubbles of vapour form within the liquid. Below this temperature evaporation takes place only from the surface.
As the temperature increases, the vapour pressure becomes high enough to form bubbles deep below the surface of the liquid (where the vapour pressure must overcome the combination of atmospheric pressure as well as the pressure from the column of liquid at depth).
This means that if one were to descend into the Challenger Deep with a heating filament, one would need to produce a very large amount of heat to create a bubble of steam. Down at the bottom of a 1 km mine shaft, the boiling point of water is 104 C°.
Conversely, as pressure decreases, the boiling point temperature also decreases. On top of Mount Everest the boiling point of water is about 65 C°.
As you travel higher, you encounter a decreasing atmospheric pressure, and so the boiling point temperature continues to decrease. This goes on until you hit the Armstrong Limit, where body fluid boiling point temperature is the same as normal core body temperature.
Water has a vapour pressure of 47mmHg at body temperature (37 C°); therefore at an atmospheric pressure of 47mmHg any exposed body fluids will boil. Those lucky few who have been exposed to such conditions and survived to report on their experience describe the peculiar sensation of their saliva boiling just before they blacked out.
Atmospheric Pressure Freezing Point (or melting point)
This is the temperature at which a substance changes state from solid to liquid. It is influenced by the solutes dissolved in the solvent: specifically, it decreases proportionally to the molar concentration of solute.
This freezing point depression is the property which governs the use of antifreeze in the cooling water of internal combustion engines. It also causes road ice to melt when salt is poured over it. Weirdly, the same concentration of solute depresses the freezing point by more degrees than it increases the boiling point.
For water, the lowest depressed freezing point seems to be around -18 C°.
Some animals have ways of making their body fluids hyperosmolar, thereby surviving temperature excesses.
Osmosis and osmotic pressure
Osmosis is the diffusion of solvent molecules into a region where there is a higher concentration of solvent.
Osmotic pressure is defined as the hydraulic pressure required to prevent the migration of solvent from the area of low solute concentration to an area of high solute concentration.
For any solution, osmotic pressure is directly proportional to its absolute temperature, and at a constant temperature, it is directly proportional to the solute concentration.
Osmotic pressure is calculated with the van 't Hoff equation:
Calculated from this equation, the total osmotic pressure of the human plasma is 5535mmHg.
This is 7.1 atmospheres.
This is the pressure one would need to apply to prevent water from moving into the plasma.
Is that a lot of pressure? The marvel and wonder of this are discussed elsewhere.
Difference between calculated and measured osmolality
The calculated osmolality of normal saline is 300, and yet its measured osmolality is 286, which resembles the measured plasma osmolality. Why, you ask? Well. The measurement of osmolality is performed with an osmometer, and it involves extrapolating osmolality on the basis of the freezing point of the liquid (see about the colligative properties of liquids). The freezing point of normal saline would suggest that its osmolality is actually 286.