What it is and when to use it
Osmotic pressure is the pressure you would have to apply to a solution to stop pure solvent flowing into it through a semipermeable membrane. It is one of the colligative properties, meaning it depends on how many dissolved particles are present rather than what they are. It is why cells swell in pure water, why IV fluids are made isotonic, and how reverse osmosis desalination must be pushed above the natural osmotic pressure of seawater.
Use this calculator for general chemistry problems, for estimating how large the driving force across a membrane is, or for checking that a lab solution is roughly matched to a target osmotic pressure. It uses the ideal, dilute-solution van 't Hoff equation, so it is most reliable for low concentrations and less reliable for concentrated or strongly interacting solutions.
The van 't Hoff equation
The calculator computes π = i × M × R × T.
- π: osmotic pressure, in atmospheres (atm).
- i: the van 't Hoff factor, the number of particles each formula unit produces. It is 1 for sugars and other non-electrolytes and ideally 2 for NaCl or 3 for CaCl₂.
- M: molarity of the solution, in mol/L.
- R: the gas constant, 0.0821 L·atm/(mol·K) in this calculator.
- T: absolute temperature in kelvin. Add 273.15 to a Celsius reading.
Worked example: 0.10 M glucose at 25 °C
Glucose does not dissociate, so i = 1. Temperature: 25 °C + 273.15 ≈ 298 K (the value entered here).
π = 1 × 0.10 × 0.0821 × 298 = 2.4466 atm, which the calculator displays as 2.447 atm. That is about 248 kPa, which is a substantial pressure for such a dilute solution.
For comparison, 0.15 M NaCl at 310 K with the ideal i = 2 gives 2 × 0.15 × 0.0821 × 310 = 7.635 atm, close to what is quoted for physiological saline.
Common mistakes and how to interpret the result
- Entering Celsius instead of kelvin. The equation needs absolute temperature, and 25 will give a result more than ten times too small.
- Forgetting i for salts. Leaving i at 1 for NaCl halves the pressure. Conversely, real electrolytes at higher concentrations show an effective i below the ideal integer value because of ion pairing.
- Mixing up molarity and molality. The formula uses molarity (mol per litre of solution); the two are nearly equal only in dilute aqueous solutions.
- Reading the result as the pressure inside the solution. It is the extra pressure that would be needed on the solution side to balance the solvent flow, not a pressure you can measure in an open beaker.