The solution certainly has properties that depend on the nature of the solute and solvent. No one would claim that a solution of arsenic is the same as a solution of sugar. However there are a few important properties of solutions that only depend on how concentrated the sample is and not on the nature of the solute. These are called Collegative Properties. There are four main Collegative Properties, vapor pressure lowering, boiling point elevation, freezing point depression and osmotic pressure. The first two of these are connected to each other so we will discuss them in this lesson. The other two will be handled in later lessons. Let's picture the surface of our solvent. There's a lot going on here. We have the jiggling and jostling of the liquid molecules. We have the vapor molecules that are exactly like the liquid ones, just with more kinetic energy. We occasionally get a liquid molecule that acquires enough kinetic energy to escape into the gas phase. And we occasionally get a vapor molecule that collides with the surface and loses enough kinetic energy to stay trapped. Evaporation and condensation in dynamic equilibrium. Think about that dynamic equilibrium. We have the rate of evaporation equaling the rate of condensation. Now how does the situation change if some of these molecules are non-volatile solute molecules? Their non-volatile means that they do not themselves vaporize. Well the rate of evaporation slows down because some of the surface molecules can't vaporize, but the rate of condensation stays the same. What will that do to the vapor pressure over the solution? The vapor pressure will decrease. Mathematically this effect is described by RyeWaltz law. Here, this symbol is the vapor pressure of the solvent when it is pure. We multiply that vapor pressure by the mole fraction of the solvent in the solution to get the vapor pressure of the solvent over the solution. What you should notice here is that this effect only depends on the concentration, not on the nature of the solute. Let's look at an example. The vapor pressure of water at 25 degrees C is 23.8 Tore. The solubility of glucose in water at that temperature is 91 grams per 100 milliliters. Let's see how we can determine the mole fraction of water in this saturated glucose solution. 91 grams of glucose is just over half a mole, and 100 milliliters of water is just over 5.5 moles. So our mole fraction of water in a saturated glucose solution is the number of moles of water over the total number of moles of water in glucose, or 0.917. And now that lets us calculate the vapor pressure of water over the solution. It is critical to recognize that in this expression for mole fraction, the denominator has to include all species in the system. So for example, if you dissolved one mole of aluminum sulfate in water, you would have 5 moles of ions. Let's think about the implications of vapor pressure lower. Recall what boiling point is. It is the temperature at which the vapor pressure of the liquid equals the surrounding atmospheric pressure. If we dissolve something in water thereby lowering the vapor pressure, we change the temperature at which the vapor pressure matches the surrounding atmospheric pressure. So the boiling point goes up. This is called boiling point elevation. If you wanted to, you could calculate the vapor pressure as a function of temperature using the Clausius Clapprion equation. And then use those results in Raoul's law to find the new boiling point, which is what I did in these graphs. But traditionally, chemists use a shortcut. We look up a boiling point elevation constant that is characteristic of the solvent we are using, multiply it by the concentration of the solute in molality, and we get the increase in the boiling point. Two details to remember. First, this concentration is in molality, not molarity. And second, it is the concentration of all solute species. So if it is an ionic compound, you have to remember to multiply by the number of ions formed from each formula unit.