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In chemistry, the osmole (Osm) is a non-SI unit of measurement that defines the number of moles of a chemical compound that contribute to a solution's osmotic pressure.

Osmolarity is a measure of the osmoles of solute per liter of solution, while the osmolality is a measure of the osmoles of solute per kilogram of solvent. Molarity and Osmolarity are not commonly used in osmometry because they are temperature dependent; that is, water changes its volume with temperature. However, if the concentration is very low, osmolarity and osmolality are considered equivalent. In calculations for these two measurements, salts are presumed to dissociate into their component ions. For example, a mole of glucose in solution is one osmole, whereas a mole of sodium chloride in solution is two osmoles (one mole of sodium and one mole of chloride). Both sodium and chloride ions affect the osmotic pressure of the solution.

The equation to determine the osmolality of a solution is given by

where

• Φ is the osmotic coefficient, which accounts for the degree of non-ideality of the solution. In the simplest case it is the degree of dissociation of the solute. Then, Φ is between 0 and 1 where 1 indicates 100% dissociation. However, Φ can also be larger than 1 (e.g. for sucrose). For salts, electrostatic effects cause Φ to be smaller than 1 even if 100% dissociation occurs (see Debye-Hückel equation).
• n is the number of particles into which a molecule dissociates. For example: Glucose equals 1 and NaCl equals 2.
• C is the molal concentration of the solution

The units are Osm/kg

Osmolality can be measured using an osmometer which measures colligative properties, such as Freezing-point depression, Vapor pressure, or Boiling-point elevation.

While similar, osmolarity and tonicity are not the same. The key difference between the two is that tonicity implies a membrane that is impermeable to the solutes on either side of it. This is not a necessary condition in the case of osmolarity. Osmolarity is a measure of the osmotically active particles in a solution and in fact makes no explicit assertion with respect to the solute permeability of any involved membranes.

The derivatives of the term: isosmotic, hyperosmotic, and hypoosmotic, should not be confused with isotonic, hypertonic and hypotonic.

Example: A urea solution that is isosmotic with respect to the cytosol of an erythrocyte is nonetheless not isotonic respective to the same erythrocyte. Urea freely diffuses across cellular membranes and is also an osmotically active particle. Normally, urea is present in a lower concentration in the nju of an erythrocyte than in an urea solution. Because urea is freely permeable to cell membranes and the concentration of urea is normally lower in the erythrocytes than in a urea solution, urea will diffuse down its concentration gradient into an erythrocyte placed into a urea solution. However, because urea is osmotically active, urea increases the solute concentration in the erythrocyte, which will then induce the osmosis of water into the cell. This can ultimately result in cell lysis. In retrospect, the isosmotic urea solution was in fact hypotonic with respect to the blood cell. Interestingly, even if the urea solution is hypoosmotic to the erythrocyte, urea will still diffuse into the cell along its concentration gradient.

• Molarity
• Molality
• Plasma osmolality
• Tonicity