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| {{About|osmolarity|the osmole unit|Osmole (unit)}}
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| '''Osmotic concentration''', formerly known as '''osmolarity''',<ref>IUPAC goldbook</ref> is the measure of [[solution|solute]] [[concentration]], defined as the number of [[osmole (unit)|osmole]]s (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 [[Mole (unit)|moles]] of solute per unit [[volume]] of solution, osmolarity measures the number of ''osmoles of solute particles'' per unit volume of solution.<ref name="Widmaier" >{{cite book | author = Widmaier, Eric P. | coauthors = Hershel Raff, Kevin T. Strang | title = Vander's Human Physiology, 11th Ed. | publisher = McGraw-Hill | pages = 108–12 | year = 2008 | isbn = 978-0-07-304962-5}}</ref>
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| Molarity and osmolarity are not commonly used in [[osmometry]] because they are [[temperature]] dependent. This is because [[H2O|water]] [[thermal expansion|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.
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| ==Types of solutes==
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| Osmolarity is distinct from molarity because it measures osmoles of solute particles rather than moles of solute. The distinction arises because some compounds can [[Dissociation (chemistry)|dissociate]] in solution, whereas others cannot.<ref name="Widmaier" />
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| [[Ionic compounds]], such as [[salt]]s, can dissociate in solution into their constituent [[ion]]s, 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<sup>+</sup> and Cl<sup>-</sup> 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.<ref name="Widmaier" />
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| 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.<ref name="Widmaier" />
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| 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.<ref name="Widmaier" />
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| ==Definition==
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| The osmolarity of a solution can be calculated from the following expression:
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| :<math> \mathrm{osmol/L} = \sum_i \varphi_i \, n_i C_i</math>
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| where
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| * ''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]]);
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| * ''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;
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| * ''C'' is the molar concentration of the solute;
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| * the index ''i'' represents the identity of a particular solute.
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| Osmolality can be measured using an [[osmometer]] which measures [[colligative properties]], such as [[Freezing-point depression]], [[Vapor pressure]], or [[Boiling-point elevation]].
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| == Osmolarity vs. tonicity ==
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| 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.<ref name="Widmaier" />
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| 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 [[Diffusion equilibrium|equilibrium]].
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| A solution can be both hyperosmotic and isotonic.<ref name="Widmaier" /> 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 ions cannot cross the membrane, it cannot draw water with it, thus causing no net change in solution volume.
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| ==Plasma osmolarity vs. osmolality==
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| Plasma osmolarity can be calculated from [[plasma osmolality]] by the following equation:<ref name=martin>Page 158 in:{{cite book |author=Martin, Alfred N.; Patrick J Sinko |title=Martin's physical pharmacy and pharmaceutical sciences: physical chemical and biopharmaceutical principles in the pharmaceutical sciences |publisher=Lippincott Williams and Wilkins |location=Phila |year=2006 |pages= |isbn=0-7817-5027-X |oclc= |doi= |accessdate=}} [http://books.google.com/books?id=nt-crAJEtVYC&printsec=frontcover#PPA158,M1]</ref>
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| Osmola'''r'''ity = osmola'''l'''ity * (ρ<sub>sol</sub> − c<sub>a</sub>)
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| Where:
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| *ρ<sub>sol</sub> is the [[density]] of the solution in g/ml, which is 1.025 g/ml for [[blood plasma]].<ref>[http://hypertextbook.com/facts/2004/MichaelShmukler.shtml Density of Blood] The Physics Factbook. Edited by Glenn Elert. Retrieved on 26 Mars, 2009</ref>
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| *c<sub>a</sub> is the ([[anhydrous]]) solute concentration in g/ml – not to be confused with the density of dried plasma
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| Since c<sub>a</sub> is slightly larger than 0.03 g/ml, plasma osmolarity is 1–2%<ref name=martin/> less than osmolality.
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| 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).
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| 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).
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| ==See also==
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| * [[Molarity]]
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| * [[Molality]]
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| * [[Plasma osmolality]]
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| * [[Tonicity]]
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| ==References==
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| D.J. Taylor, N.P.O. Green, G.W. Stout Biological Science
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| {{Reflist}}
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| ==External links==
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| [http://goldbook.iupac.org/O04343.html]
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| [[Category:Analytical chemistry]]
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| [[Category:Amount of substance]]
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| [[Category:Solutions]]
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