Osmosis and osmotic pressure

Semipermeable membranes, natural (pig's bladder, parchment) or synthetic (cellophane), have submicroscopic pores that let small solvent molecules such as water through but hinder larger solute particles.

Osmosis is the net flow of solvent through a semipermeable membrane from pure solvent, or a dilute solution, into a more concentrated solution.

Osmotic pressure is the excess pressure that must be applied to the solution side to just stop osmosis. It is a colligative property.

For dilute solutions π = CRT, where C is molarity; with n₂ mol of solute in V litres, π = (n₂/V)RT, and with w₂ g of solute, M₂ = w₂RT/(πV).

Osmotic pressure is the method of choice for molar masses of proteins, polymers and other macromolecules: it is measured near room temperature, uses molarity, and gives a sizeable reading even for very dilute solutions. Biomolecules are often unstable when heated, and polymers are poorly soluble.

Solutions with equal osmotic pressure at a temperature are isotonic, and no osmosis occurs between them. The fluid in blood cells is isotonic with 0.9% (mass/volume) sodium chloride, normal saline, which is safe for intravenous injection.

Above 0.9% NaCl (hypertonic) water leaves blood cells and they shrink; below 0.9% (hypotonic) water enters and they swell.

Everyday osmosis: raw mangoes shrivel in brine, wilted flowers and limp carrots firm up in fresh water, salty diets cause water retention (edema), water rises into plants partly by osmosis, and salted meat or sugared fruit kills bacteria by drawing water out of them.

Reverse osmosis: a pressure larger than the osmotic pressure applied to the solution side pushes pure solvent out through the membrane. It is used to desalinate sea water, often with a cellulose acetate film that passes water but not ions and impurities.