Solutes, such as proteins or simple ions, dissolve in a solvent such as water. This raises the concentration of the solute in these areas. The solvent then diffuses to these areas of higher solute concentration to equalize the concentration of the solute throughout the solution.
Example of osmosis
A practical example of this osmosis in cells can be seen in red blood cells. These contain a high concentration of solutes including salts and protein. When the cells are placed in solution, water rushes in to the area of high solute concentration, bursting the cell.
Many plant cells do not burst in the same experiment. This is because the osmotic entry of water is opposed and eventually equalled by the pressure exerted by the cell wall, creating a steady state[?]. In fact, osmotic pressure is the main cause of support in plant leaves.
When a plant cell is placed in a solution higher in solutes than inside the cell osmosis out of the cell occurs. The water in the cell moves to an area higher in solute concentration, and the cell shrinks and so becomes flaccid. This means the cell has become plasmolysed - the cell membrane has completely left the cell wall due to lack of water pressure on it.
When a solute is dissolved in a solvent, the random mixing of the two species results in an increase in the entropy of the system, which corresponds to a reduction in the chemical potential. For the case of an ideal solution[?] the reduction in chemical potential corresponds to:
where R is the ideal gas constant, T is the temperature and x2 is the solute concentration in terms of mole fraction. Most real solutions approximate ideal behavior for low solute concentrations (At higher concentrations interactions between solute and solvent cause deviations from Equation 1). This reduced potential creates a driving force and it is this force which drives diffusion of water through the semipermeable membrane.
As mentioned before, osmosis can be opposed by increasing the pressure in the region of high solute concentration with respect to that in the low solute concentration region. The pressure differential at which the flow of solvent through the membrane is stopped is called the osmotic pressure or turgor.
Increasing the pressure increases the chemical potential of the system in proportion to the molar volume (δμ = δPV). Therefore, osmosis stops, when the increase in potential due to pressure equals the potential decrease from Equation 1, i.e.:
Where δP is the osmotic pressure and V is the molar volume of the solvent.
For the case of very low solute concentrations, ln(1-x2) ~ x2 and Equation 2 can be rearranged into the following expression for osmotic pressure:
The osmosis process can be driven in reverse with solvent moving from a region of high solute concentration to a region of low solute concentration by applying a pressure in excess of the osmotic pressure. This reverse osmosis technique is commonly applied to purify water.