Is osmosis really a special case of diffusion across a semipermeable membrane?

Question

A general idea we have about osmosis is that it is the movement of solvent particles through a semi permeable membrane from a region of high concentration to a region of low concentration. This seems correct in case of a gas. Gaseous particles spread through any container in which they are kept, which makes it seem to us as if they moved from high concentration to lower concentration. But in the case of a liquid how is this true? For example consider a bottle half filled with water, water never moves from high concentration to lower concentration inside the bottle.

So is it wrong to explain osmosis as a special case of diffusion?

What actually makes solvent particles move from regions of higher concentration to regions of lower concentration? Is it solute attraction? Or a difference in pressure in pure solvent phase and solution phase?

Answer 1

I've listed a series of points about diffusion and Osmosis.

The cause of osmosis is simply diffusion: the solvent can diffuse through the membrane and the solute cannot.

Thermodynamically, we say that the reason that the diffusion occurs is that the chemical potential of the solution is at a lower chemical potential than the pure solvent and so there is a tendency for the solvent to pass into its solution and to dilute it. The osmotic pressure is the excess pressure just needed to stop that flow. If more pressure were to be applied the flow of solvent would be reversed and the solution would become more concentrated.

Osmosis occurs because there is an increase in entropy overall when the solvent dilutes the solution. There is a decrease in entropy in the pure solvent side because its volume decreases but this is balanced by an increase in entropy on the solution side as the volume increases so this is not the reason osmosis occurs. The reason is the increase in entropy that is experienced by the solute now in a larger volume or equivalently more solvent molecules in which to dissolve. This means that there are more ways of arranging solute and solvent molecules compared to that before osmosis stared, and this represents an increase in entropy. (Left unchecked this will continue until the solution is so dilute it is indistinguishable from the solvent).

When we say that a solution has an osmotic pressure of, say, 5 atm, it does not mean that the solution exerts this pressure but that if it is in contact with its solvent through a semi-permeable membrane then an applied pressure of 5 atm is needed to maintain equilibrium.

It is possible for a solute in a solution to diffuse spontaneously from regions of higher to lower concentration so that uniform concentration is reached and the solution is at equilibrium. Naturally, diffusion does not stop at this point but is always present both for solute and solvent.

In some cases, however, diffusion can occur form lower to higher concentrations. Such a situation is that when a solute is dissolved in two immiscible solvents in contact with one another, and in which the solute has different solubilities. Solute diffusion will occur until the chemical potential in each solvent is equal which can mean diffusion into the more concentrated solution. Thus at equilibrium there will be a large difference in solute concentration but no overall flow of solute due to diffusion.

Answer 2

A general idea we have about osmosis is that it is the movement of solvent particles through a semi permeable membrane from a region of high concentration to a region of low concentration.

You have to carefully specify which concentration you refer to. The net transport is from regions of high solvent concentration (low solute concentration) to low solvent concentration (high solute concentration).

This seems correct in case of a gas. Gaseous particles spread through any container in which they are kept, which makes it seem to us as if they moved from high concentration to lower concentration.

It does not just seem as if, that is what actually happens as net effect. The reason that more particles move out of a region of high concentration is because there is more of them.

But in the case of a liquid how is this true? For example consider a bottle half filled with water, water never moves from high concentration to lower concentration inside the bottle.

That would be an example of a phase change. Water molecules do move from the liquid to the gas phase and vice versa. However, because the water molecules interact in the liquid and don't interact in the gas phase, most water molecules are in the liquid phase. The reason that some are in the gas phase is that they are more dispersed in that phase and mix with other components of air, both of which is entropically favored.

So is it wrong to explain osmosis as a special case of diffusion?

No, it is perfectly alright to explain osmosis with diffusion. There is a higher concentration of solvent on the side of the membrane with lower solute concentrations, so more solvent molecules move away from that side than move to that side.

What actually makes solvent particles move from regions of higher concentration to regions of lower concentration? Is it solute attraction?

No, this also works for ideal solutions, where the solute:solvent interactions are very similar to the solvent:solvent interactions. It is a process driven by entropy, not enthalpy.

Or a difference in pressure in pure solvent phase and solution phase?

In a system where the liquids have room to expand, the pressure is the same. In a system where the liquid compartments are kept at constant volume, the pressure on the side of higher solute concentration increases (because the amount of solvent in that compartment increases through osmosis). So the pressure difference is caused by the osmosis, not the other way around.

There is a process called reverse osmosis used to purify water, where the pressure differential results in water going from the "dirty" compartment to the "clean" compartment.