When the temperature of a liquid is below its boiling point, we can assume that the only molecules that can escape from the liquid to form a gas are those that lie near the surface of the liquid. When a solute is added to the solvent, some of the solute molecules occupy the space near the surface of the liquid, as shown in the figure below. When a solute is dissolved in a solvent, the number of solvent molecules near the surface decreases, and the vapor pressure of the solvent decreases.
This has no effect on the rate at which solvent molecules in the gas phase condense to form a liquid. But it decreases the rate at which the solvent molecules in the liquid can escape into the gas phase.
As a result, the vapor pressure of the solvent escaping from a solution should be smaller than the vapor pressure of the pure solvent. Between and , Francois-Marie Raoult showed that the vapor pressure of a solution is equal to the mole fraction of the solvent times the vapor pressure of the pure liquid. This equation, which is known as Raoult's law , is easy to understand. When the solvent is pure, and the mole fraction of the solvent is equal to 1, P is equal to P o. As the mole fraction of the solvent becomes smaller, the vapor pressure of the solvent escaping from the solution also becomes smaller.
Let's assume, for the moment, that the solvent is the only component of the solution that is volatile enough to have a measurable vapor pressure.
If this is true, the vapor pressure of the solution will be equal to the vapor pressure of the solvent escaping from the solution. Raoult's law suggests that the difference between the vapor pressure of the pure solvent and the solution increases as the mole fraction of the solvent decreases.
The change in the vapor pressure that occurs when a solute is added to a solvent is therefore a colligative property. If it depends on the mole fraction of the solute, then it must depend on the ratio of the number of particles of solute to solvent in the solution but not the identity of the solute.
The figure below shows the consequences of the fact that solutes lower the vapor pressure of a solvent. The solid line connecting points B and C in this phase diagram contains the combinations of temperature and pressure at which the pure solvent and its vapor are in equilibrium. Each point on this line therefore describes the vapor pressure of the pure solvent at that temperature. The dotted line in this figure describes the properties of a solution obtained by dissolving a solute in the solvent.
At any given temperature, the vapor pressure of the solvent escaping from the solution is smaller than the vapor pressure of the pure solvent. The dotted line therefore lies below the solid line. According to this figure, the solution can't boil at the same temperature as the pure solvent. If the vapor pressure of the solvent escaping from the solution is smaller than the vapor pressure of the pure solvent at any given temperature, the solution must be heated to a higher temperature before it boils.
The lowering of the vapor pressure of the solvent that occurs when it is used to form a solution therefore increases the boiling point of the liquid. When phase diagrams were introduced, the triple point was defined as the only combination of temperature and pressure at which the gas, liquid, and solid can exist at the same time. When a solute is dissolved in a solvent, the boiling point of the solution is raised according to the equation:.
Want to practice some calculations using this equation? Click here. Let's start this discussion the same way we started the others, by defining the normal freezing point. The normal freezing point of a liquid is is the temperature at which a liquid becomes a solid at 1 atm. A more specific definition of freezing point is the temperature at which solid and liquid phases coexist in equilibrium.
Let's see if we can figure out why the freezing point is lowered when we add solutes to a solution. We already know that in order to freeze a liquid, we have to lower the temperature.
As the temperature lowers, the solution becomes more ordered as it moves toward the solid phase. This is an effect that works against the second law of thermodynamics. In short, entropy disorder likes to increase not decrease in the natural scheme of things. So if we have to lower the temperature to a certain point to freeze a pure solvent, when we add a solute we add to the entropy of the system, right? The mixture is more disordered than the pure. This additional amount of entropy must now be overcome to allow the liquid to change phases into a solid become ordered.
This means that the temperature will have to be even lower than before. While the chemical nature of the solute is not a factor, it is necessary to take into account whether the solute is an electrolyte or a nonelectrolyte. Recall that ionic compounds are strong electrolytes and thus dissociate into ions when they dissolve. This results in a larger number of dissolved particles.
For example, consider two different solutions of equal concentration. One is made from the ionic compound sodium chloride, while the other is made from the molecular compound glucose. The following equations show what happens when these solutes dissolve. The sodium chloride dissociates into two ions, while the glucose does not dissociate. Therefore, equal concentrations of each solution will result in twice as many dissolved particles in the case of the sodium chloride.
The vapor pressure of the sodium chloride solution will be lowered twice the amount as the glucose solution. Watch the video at the link given and answer the following questions:.
Skip to main content. Previous Unit Kinetic-Molecular Theory. Next Unit Chemical Reaction Rates. Kendal Orenstein. Thank you for watching the video. Start Your Free Trial Learn more. Kendal Orenstein Rutger's University M. Explanation Transcript Vapor pressure lowering is a colligative property of solutions. Chemistry Chemical Solutions. Science Biology Chemistry Physics.
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