Osmolarity

Osmotic concentration, formerly known as osmolarity, is the measure of solute concentration, defined as the number of osmoles (Osm) of solute per litre (L) of solution (osmol/L or Osm/L). The osmolarity of a solution is usually expressed as Osm/L (pronounced "osmolar"), in the same way that the molarity of a solution is expressed as "M" (pronounced "molar"). Whereas molarity measures the number of moles of solute per unit volume of solution, osmolarity measures the number of osmoles of solute particles per unit volume of solution. This value allows the measurement of the osmotic pressure of a solution and the determination of how the solvent will diffuse across a semipermeable membrane (osmosis) separating two solutions of different osmotic concentration.

Osmolarity is distinct from molarity because it measures osmoles of solute particles rather than moles of solute. The distinction arises because some compounds can dissociate in solution, whereas others cannot.

Ionic compounds, such as salts, can dissociate in solution into their constituent ions, so there is not a one-to-one relationship between the molarity and the osmolarity of a solution. For example, sodium chloride (NaCl) dissociates into Na⁺ and Cl⁻ ions. Thus, for every 1 mole of NaCl in solution, there are 2 osmoles of solute particles (i.e., a 1 mol/L NaCl solution is a 2 osmol/L NaCl solution). Both sodium and chloride ions affect the osmotic pressure of the solution.

Another example is magnesium chloride (MgCl2), which dissociates into Mg2+ and 2Cl- ions. For every 1 mole of MgCl2 in the solution, there are 3 osmoles of solute particles.

Nonionic compounds do not dissociate, and form only 1 osmole of solute per 1 mole of solute. For example, a 1 mol/L solution of glucose is 1 osmol/L.

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