osmotic pressure
/ oz-MOT-ik /
Put pure water on one side of a thin membrane and salty water on the other, where the membrane lets water through but not the salt. Water seeps across toward the salty side, as if trying to dilute it, and the salty side swells and rises. Osmotic pressure is the push you would need to apply to the salty side to stop that flow exactly.
Osmosis is water moving across a selective membrane from where it is more pure (low solute) to where it is less pure (high solute). The osmotic pressure is the pressure that just halts this movement; the more crowded the solute, the harder water tries to come in, and the higher that pressure. For dilute solutions it follows a strikingly gas-like law: pressure rises in step with the concentration of solute particles and with temperature.
This quiet pressure runs much of biology. It is how plant roots draw water up and how cells stay plump or shrivel; medical fluids must be matched to blood's osmotic pressure or red cells will burst or collapse. Reverse osmosis turns the idea around — push hard enough on salty water and you force pure water back out through the membrane, which is how seawater is desalinated.
Soak a raisin in water and it plumps up: water flows in through the skin toward the sugary inside. Soak a grape in syrup and it shrivels: water flows out toward the sweeter outside.
Water crosses a membrane toward the more concentrated side; osmotic pressure is what stops it.
Osmotic pressure is a colligative property, so it counts particles: a solute that breaks into ions exerts a proportionally larger osmotic pressure. Even quite dilute solutions can develop surprisingly large osmotic pressures, which makes the effect sensitive enough to weigh huge molecules like proteins.