## Osmolarity and Osmolality

Another unit of measurement is osmolality, which determines the distribution of water among the different fluid compartments, particularly between the extracellular and intracellular fluids. Osmolality effects this distribution of water through the generation of osmotic pressure.

The osmotic pressure generated by a solution is proportional to the number of particles per unit volume of solvent, not to the type, valence, or weight of the particles.

The unit of measurement of osmolality is the osmole. The term relates to osmosis and is typically used for osmotically active solutions. The osmolality of a solution is the number of osmoles of solute per kilogram of solvent. One osmole (Osmol) is defined as one gram molecular weight (1 mole) of any non-dissociable substance (such as glucose) that contains 6.02 x 1023 particles and contributes to a solution’s osmotic pressure. In the relatively dilute fluids in the body, the osmolality is measured in milliosmols (one-thousandth of an osmole) per kilogram of water (mOsmol/kg). Osmolarity is similar but is defined as the number of osmoles (or mOsmol) per liter of solvent. Since most solutes are measured in the laboratory in units of millimoles per liter, milligrams per deciliter, or milliequivalents per liter, the following formulae must be used to convert into mOsmol/kg:

mOsmol/kg   =     n  x   mmol/L

mOsmol/kg    =    (n   x   mg/dL  x  10)   ÷   mol wt

mOsmol/kg    =    (n   x   mEq/L)   ÷   valence

where n is the number of dissociable particles per molecule. When n = 1, as for Na+, Cl, Ca2+, urea, and glucose, 1 mmol/L equals 1 mOsmol/kg. If, however, a compound dissociates into two or more particles, 1 mmol/L will generate an osmotic effect greater than 1 mOsmol/kg. For example, a solution of 1 mol/L NaCl corresponds to an osmolarity of 2 osmol/L. The NaCl salt particle dissociates fully in water to become two separate particles: a Na+ ion and a Cl ion. Therefore, each mole of NaCl becomes two osmoles in solution, one mole of Na+ and one mole of Cl(1).

In the laboratory, the osmotic concentration of a solution is measured not as an osmotic pressure but according to other properties of solutions (known as colligative properties) such as their ability to depress the freezing point or elevate the boiling point of water. Solute-free water freezes at 0° C. If 1 Osmol of any solute (or combination of solutes) is added to 1 kg of water, the freezing point of this water will be depressed by 1.86° C. This observation can be used to calculate the osmotic concentration of a solution. For example, the freezing point of the plasma water is normally about -0.521° C. This represents an osmolality of 0.280 Osmol/kg (0.521/1.86) or 280 mOsmol/kg.

Importantly, only solutes that cannot cross the semipermeable membrane separating two compartments generate an effective osmotic pressure. Thus, urea, which can cross the cell membrane, does not contribute to osmotic pressure but will be measured as part of the plasma osmolality by freezing point depression. There is therefore a difference between the total osmolality and the effective osmolality of a solution, with the latter being determined only by osmotically active solutes (such as Na+ and K+ across the cell membrane).

#### References:

1.    Scudiero, Louis (2006), Molecular Weight Determination of Sucrose Using an Osmometer, Washington State University

P/N 101833-01 Rev A 07/2012