Why is osmotic pressure a colligative property

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. The figure above shows that the triple point of the solution occurs at a lower temperature than the triple point of the pure solvent. By itself, the change in the triple point is not important. The van 't Hoff model imagines that a substance dissolved in a fluid medium behaves as if it were a gas in a vacuum.

In short, the salt in the beaker of water exerts the same pressure on the walls of the beaker as it would if the water was magically removed, leaving the salt in a weird gaseous form.

Turns out, this works, or at least as long as the solutes are reasonably dilute. Other variations exist. For example, if you assume the solution is behaving ideally, you can simplify this even more, as:. At this stage it would be possible to launch into pages of derivations, but it would not be useful, as it would not add very much to the otherwise rather basic points this chapter is trying to make.

The bottom line is that there is a considerable amount of pressure being applied to fluids because of osmotic gradients. The same value can be seen in other reputable resources. Also, speculating a little about this concept: if pressure can oppose the movement of the solvent into an area of higher solute concentration, then, surely, with enough pressure Yes, the application of hydraulic pressure to the solution can be used to force solvent through the semipermeable membrane.

The pressures used in commercially available personal reverse osmosis water purifiers is usually about 2. The concentrated effluent is discarded. This is usually how SLED dialysis units are operated, as it would not be possible to supply enough premixed dialysate fluid with the sort of flow rate they require.

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CBSE has released the term-1 admit card Share This Video. Apne doubts clear karein ab Whatsapp par bhi. Try it now. Note that the osmotic pressure is the pressure required to stop osmosis, not to sustain it.

It is important to understand that this means nothing more than that a pressure of this value must be applied to the solution to prevent flow of pure solvent into this solution through a semipermeable membrane separating the two liquids. The Dutch scientist Jacobus Van't Hoff was one of the giants of physical chemistry.

In contrast to the need to employ solute molality to calculate the effects of a non-volatile solute on changes in the freezing and boiling points of a solution, we can use solute molarity to calculate osmotic pressures.

In this context, molarity refers to the summed total of the concentrations of all solute species. As such, this equation gives valid results only for extremely dilute "ideal" solutions. Sea water contains dissolved salts at a total ionic concentration of about 1. What pressure must be applied to prevent osmotic flow of pure water into sea water through a membrane permeable only to water molecules? Since all of the colligative properties of solutions depend on the concentration of the solvent, their measurement can serve as a convenient experimental tool for determining the concentration, and thus the molecular weight, of a solute.

Osmotic pressure is especially useful in this regard, because a small amount of solute will produce a much larger change in this quantity than in the boiling point, freezing point, or vapor pressure. Molecular weight determinations are very frequently made on proteins or other high molecular weight polymers.

These substances, owing to their large molecular size, tend to be only sparingly soluble in most solvents, so measurement of osmotic pressure is often the only practical way of determining their molecular weights. The osmotic pressure of a benzene solution containing 5. Estimate the average molecular weight of the polystyrene in this sample.