WordNet

(of a solution) the extent to which a solution has a higher osmotic pressure than some other
(of living tissue) in a state of abnormally high tension; "hypertonic muscle tissue"
(of a solution) having a higher osmotic pressure than a comparison solution

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出典(authority):フリー百科事典『ウィキペディア（Wikipedia）』「2013/02/05 15:32:12」(JST)

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This article is about osmolarity. For the osmole unit, see Osmole (unit).

Osmotic concentration, formerly known as osmolarity,^[1] 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.^[2]

Osmolality is a measure of the osmoles of solute per kilogram of solvent (osmol/kg or Osm/kg).

Molarity and osmolarity are not commonly used in osmometry because they are temperature dependent. This is because water changes its volume with temperature (See: Vapour pressure of water). However, if the concentration of solutes is very low, osmolarity and osmolality are considered equivalent.

Types of solutes

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.^[2]

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 M NaCl solution is a 2 Osm NaCl solution). Both sodium and chloride ions affect the osmotic pressure of the solution.^[2]

Nonionic compounds do not dissociate, and form only 1 osmole of solute per 1 mole of solute. For example, a 1 M solution of glucose is 1 Osm.^[2]

Multiple compounds may contribute to the osmolarity of a solution. For example, a 3 Osm solution might consist of: 3 moles glucose, or 1.5 moles NaCl, or 1 mole glucose + 1 mole NaCl, or 2 moles glucose + 0.5 mole NaCl, or any other such combination.^[2]

Definition

The osmolarity of a solution can be calculated from the following expression:

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 (e.g. ions) into which a molecule dissociates. For example: glucose has n of 1, while NaCl has n of 2;
C is the molar concentration of the solute;
the index i represents the identity of a particular solute.

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

Osmolarity vs. tonicity

Osmolarity and tonicity are related, but different concepts. Thus, the terms ending in -osmotic (isosmotic, hyperosmotic, hyposmotic) are not synonymous with the terms ending in -tonic (isotonic, hypertonic, hypotonic). The terms are related in that they both compare the solute concentrations of two solutions separated by a membrane. The terms are different because osmolarity takes into account the total concentration of penetrating solutes and non-penetrating solutes, whereas tonicity takes into account the total concentration of only non-penetrating solutes.^[2]

Penetrating solutes can diffuse through the cell membrane, causing momentary changes in cell volume as the solutes "pull" water molecules with them. Non-penetrating solutes cannot cross the cell membrane, and therefore osmosis of water must occur for the solutions to reach equilibrium.

A solution can be both hyperosmotic and isotonic.^[2] For example, the intracellular fluid and extracellular can be hyperosmotic, but isotonic – if the total concentration of solutes in one compartment is different from that of the other, but one of the ions can cross the membrane, drawing water with it and thus causing no net change in solution volume.

Plasma osmolarity vs. osmolality

Plasma osmolarity can be calculated from plasma osmolality by the following equation:^[3]

Osmolarity = osmolality * (ρ_sol − c_a)

Where:

ρ_sol is the density of the solution in g/ml, which is 1.025 g/ml for blood plasma.^[4]
c_a is the (anhydrous) solute concentration in g/ml – not to be confused with the density of dried plasma

Since c_a is slightly larger than 0.03 g/ml, plasma osmolarity is 1–2%^[3] less than osmolality.

According to IUPAC, osmolality is the quotient of the negative natural logarithm of the rational activity of water and the molar mass of water, whereas osmolarity is the product of the osmolality and the mass density of water (also known as osmotic concentration).

In simpler terms, osmolality is an expression of solute osmotic concentration per mass, whereas osmolarity is per volume of solvent (thus the conversion by multiplying with the mass density).

References

D.J. Taylor, N.P.O. Green, G.W. Stout Biological Science

^ IUPAC goldbook
^ ^a ^b ^c ^d ^e ^f ^g Widmaier, Eric P.; Hershel Raff, Kevin T. Strang (2008). Vander's Human Physiology, 11th Ed.. McGraw-Hill. pp. 108–12. ISBN 978-0-07-304962-5.
^ ^a ^b Page 158 in:Martin, Alfred N.; Patrick J Sinko (2006). Martin's physical pharmacy and pharmaceutical sciences: physical chemical and biopharmaceutical principles in the pharmaceutical sciences. Phila: Lippincott Williams and Wilkins. ISBN 0-7817-5027-X. [1]
^ Density of Blood The Physics Factbook. Edited by Glenn Elert. Retrieved on 26 Mars, 2009

External links

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