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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. One osmole (Osmol) is defined as one gram molecular weight (1 mole) of any non-dissociable substance (such as glucose) and contains 6.02 x 1023 particles. 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 measured as osmoles (or mOsmol) per liter of solvent. Since most solutes are measured in the laboratory in units of millimols 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, at the concentrations present in the body, ionic interactions reduce the random movement of NaCl so that it acts is if it were only 75 percent (not 100 percent) dissociated. Thus, for each 1 mmol/l of NaCl, there will be 0.75 mmol/L each of Na+ and Cl- and 0.25 mmol/L of NaCl or 1.75 mOsmol/kg (2 x 0.75 = 1.5 + 0.25 = 1.75).
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 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.
However, only solutes that cannot cross the 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).
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