Osmotic concentration

Last updated March 23, 2023

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] This value allows the measurement of the osmotic pressure of a solution and the determination of how the solvent will diffuse across a semipermeable membrane (osmosis) separating two solutions of different osmotic concentration.

Unit

The unit of osmotic concentration is the osmole. This is a non-SI unit of measurement that defines the number of moles of solute that contribute to the osmotic pressure of a solution. A milliosmole (mOsm) is 1/1,000 of an osmole. A microosmole (μOsm) (also spelled micro-osmole) is 1/1,000,000 of an osmole.

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 mol/L NaCl solution is a 2 osmol/L NaCl solution). Both sodium and chloride ions affect the osmotic pressure of the solution.^[2]

Another example is magnesium chloride (MgCl₂), which dissociates into Mg²⁺ and 2Cl⁻ ions. For every 1 mole of MgCl₂ in the solution, there are 3 osmoles of solute particles.

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.^[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, given in osmoles per liter (osmol/L) is calculated from the following expression:

\mathrm {osmolarity} =\sum _{i}\varphi _{i}\,n_{i}C_{i}

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.

Osmolarity 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 distinct 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 non-freely penetrating solutes only.^[3]^[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; therefore, the movement of water across the cell membrane (i.e., osmosis) 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 (in other words, a penetrating solute), drawing water with it, thus causing no net change in solution volume.

Plasma osmolarity vs. osmolality

Plasma osmolarity can be calculated from plasma osmolality by the following equation:^[4]

Osmolarity = osmolality

\times (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.^[5]
$c a$ is the (anhydrous) solute concentration in g/ml – not to be confused with the density of dried plasma

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).^[1]

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

{\text{osmolality}}=\sum _{i}\varphi _{i}\,n_{i}m_{i}

where $m i$ is the molality of component $i$ .

Plasma osmolarity/osmolality is important for keeping proper electrolytic balance in the blood stream. Improper balance can lead to dehydration, alkalosis, acidosis or other life-threatening changes. Antidiuretic hormone (vasopressin) is partly responsible for this process by controlling the amount of water the body retains from the kidney when filtering the blood stream.^[6]

Related Research Articles

Osmotic pressure is the minimum pressure which needs to be applied to a solution to prevent the inward flow of its pure solvent across a semipermeable membrane. It is also defined as the measure of the tendency of a solution to take in its pure solvent by osmosis. Potential osmotic pressure is the maximum osmotic pressure that could develop in a solution if it were separated from its pure solvent by a semipermeable membrane.

Solubility equilibrium is a type of dynamic equilibrium that exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The solid may dissolve unchanged, with dissociation, or with chemical reaction with another constituent of the solution, such as acid or alkali. Each solubility equilibrium is characterized by a temperature-dependent solubility product which functions like an equilibrium constant. Solubility equilibria are important in pharmaceutical, environmental and many other scenarios.

Molar concentration is a measure of the concentration of a chemical species, in particular of a solute in a solution, in terms of amount of substance per unit volume of solution. In chemistry, the most commonly used unit for molarity is the number of moles per liter, having the unit symbol mol/L or mol/dm³ in SI unit. A solution with a concentration of 1 mol/L is said to be 1 molar, commonly designated as 1 M.

Passive transport is a type of membrane transport that does not require energy to move substances across cell membranes. Instead of using cellular energy, like active transport, passive transport relies on the second law of thermodynamics to drive the movement of substances across cell membranes. Fundamentally, substances follow Fick's first law, and move from an area of high concentration to one of low concentration because this movement increases the entropy of the overall system. The rate of passive transport depends on the permeability of the cell membrane, which, in turn, depends on the organization and characteristics of the membrane lipids and proteins. The four main kinds of passive transport are simple diffusion, facilitated diffusion, filtration, and/or osmosis.

In chemistry, colligative properties are those properties of solutions that depend on the ratio of the number of solute particles to the number of solvent particles in a solution, and not on the nature of the chemical species present. The number ratio can be related to the various units for concentration of a solution such as molarity, molality, normality (chemistry), etc. The assumption that solution properties are independent of nature of solute particles is exact only for ideal solutions, which are solutions that exhibit thermodynamic properties analogous to those of an ideal gas, and is approximate for dilute real solutions. In other words, colligative properties are a set of solution properties that can be reasonably approximated by the assumption that the solution is ideal.

<span class="mw-page-title-main">Loop of Henle</span> Part of kidney tissue

In the kidney, the loop of Henle is the portion of a nephron that leads from the proximal convoluted tubule to the distal convoluted tubule. Named after its discoverer, the German anatomist Friedrich Gustav Jakob Henle, the loop of Henle's main function is to create a concentration gradient in the medulla of the kidney.

The van 't Hoff factor $i$ is a measure of the effect of a solute on colligative properties such as osmotic pressure, relative lowering in vapor pressure, boiling-point elevation and freezing-point depression. The van 't Hoff factor is the ratio between the actual concentration of particles produced when the substance is dissolved and the concentration of a substance as calculated from its mass. For most non-electrolytes dissolved in water, the van 't Hoff factor is essentially 1. For most ionic compounds dissolved in water, the van 't Hoff factor is equal to the number of discrete ions in a formula unit of the substance. This is true for ideal solutions only, as occasionally ion pairing occurs in solution. At a given instant a small percentage of the ions are paired and count as a single particle. Ion pairing occurs to some extent in all electrolyte solutions. This causes the measured van 't Hoff factor to be less than that predicted in an ideal solution. The deviation for the van 't Hoff factor tends to be greatest where the ions have multiple charges.

Saline is a mixture of sodium chloride (salt) and water. It has a number of uses in medicine including cleaning wounds, removal and storage of contact lenses, and help with dry eyes. By injection into a vein it is used to treat dehydration such as that from gastroenteritis and diabetic ketoacidosis. Large amounts may result in fluid overload, swelling, acidosis, and high blood sodium. In those with long-standing low blood sodium, excessive use may result in osmotic demyelination syndrome.

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The ionic strength of a solution is a measure of the concentration of ions in that solution. Ionic compounds, when dissolved in water, dissociate into ions. The total electrolyte concentration in solution will affect important properties such as the dissociation constant or the solubility of different salts. One of the main characteristics of a solution with dissolved ions is the ionic strength. Ionic strength can be molar or molal and to avoid confusion the units should be stated explicitly. The concept of ionic strength was first introduced by Lewis and Randall in 1921 while describing the activity coefficients of strong electrolytes.

Osmol gap in medical science is the difference between measured serum osmolality and calculated serum osmolality.

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Osmosis is the spontaneous net movement or diffusion of solvent molecules through a selectively-permeable membrane from a region of high water potential to a region of low water potential, in the direction that tends to equalize the solute concentrations on the two sides. It may also be used to describe a physical process in which any solvent moves across a selectively permeable membrane separating two solutions of different concentrations. Osmosis can be made to do work. Osmotic pressure is defined as the external pressure required to be applied so that there is no net movement of solvent across the membrane. Osmotic pressure is a colligative property, meaning that the osmotic pressure depends on the molar concentration of the solute but not on its identity.

In thermodynamics, an apparent molar property of a solution component in a mixture or solution is a quantity defined with the purpose of isolating the contribution of each component to the non-ideality of the mixture. It shows the change in the corresponding solution property per mole of that component added, when all of that component is added to the solution. It is described as apparent because it appears to represent the molar property of that component in solution, provided that the properties of the other solution components are assumed to remain constant during the addition. However this assumption is often not justified, since the values of apparent molar properties of a component may be quite different from its molar properties in the pure state.

Stool osmotic gap is a measurement of the difference in solute types between serum and feces, used to distinguish among different causes of diarrhea.

A balanced salt solution (BSS) is a solution made to a physiological pH and isotonic salt concentration. Solutions most commonly include sodium, potassium, calcium, magnesium, and chloride. Balanced salt solutions are used for washing tissues and cells and are usually combined with other agents to treat the tissues and cells. They provide the cells with water and inorganic ions, while maintaining a physiological pH and osmotic pressure.

The rock dove, Columbia livia, has a number of special adaptations for regulating water uptake and loss.

References

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

1 2 McNaught, A. D.; Wilkinson, A.; Chalk, S. J. (1997). IUPAC. Compendium of Chemical Terminology (the "Gold Book") (2nd ed.). Oxford: Blackwell Scientific Publications. ISBN 0-9678550-9-8 . Retrieved 23 January 2022.
1 2 3 4 5 6 7 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.
↑ Costanzo, Linda S. (2017-03-15). Physiology. Preceded by: Costanzo, Linda S., 1947- (Sixth ed.). Philadelphia, PA. ISBN 9780323511896. OCLC 965761862.
↑ Martin, Alfred N.; Patrick J Sinko (2006). Martin's physical pharmacy and pharmaceutical sciences: physical chemical and biopharmaceutical principles in the pharmaceutical sciences. Philadelphia, Pennsylvania: Lippincott Williams and Wilkins. p. 158. ISBN 0-7817-5027-X.
↑ Shmukler, Michael (2004). Elert, Glenn (ed.). "Density of blood". The Physics Factbook. Retrieved 2022-01-23.
↑ Earley, L. E.; Sanders, C. A. (1959). "The Effect of Changing Serum Osmolality on the Release of Antidiuretic Hormone in Certain PAtients with Decompensated Cirrhosis of the Liver and Low Serum Osmolality". Journal of Clinical Investigation. 38 (3): 545–550. doi:10.1172/jci103832. PMC 293190 . PMID 13641405.

External links

Online Serum Osmolarity/Osmolality calculator

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[gold_book-1] 1 2 McNaught, A. D.; Wilkinson, A.; Chalk, S. J. (1997). IUPAC. Compendium of Chemical Terminology (the "Gold Book") (2nd ed.). Oxford: Blackwell Scientific Publications. ISBN 0-9678550-9-8 . Retrieved 23 January 2022.

[Widmaier-2] 1 2 3 4 5 6 7 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.

[3] Costanzo, Linda S. (2017-03-15). Physiology. Preceded by: Costanzo, Linda S., 1947- (Sixth ed.). Philadelphia, PA. ISBN 9780323511896. OCLC 965761862.

[martin-4] Martin, Alfred N.; Patrick J Sinko (2006). Martin's physical pharmacy and pharmaceutical sciences: physical chemical and biopharmaceutical principles in the pharmaceutical sciences. Philadelphia, Pennsylvania: Lippincott Williams and Wilkins. p. 158. ISBN 0-7817-5027-X.

[5] Shmukler, Michael (2004). Elert, Glenn (ed.). "Density of blood". The Physics Factbook. Retrieved 2022-01-23.

[6] Earley, L. E.; Sanders, C. A. (1959). "The Effect of Changing Serum Osmolality on the Release of Antidiuretic Hormone in Certain PAtients with Decompensated Cirrhosis of the Liver and Low Serum Osmolality". Journal of Clinical Investigation. 38 (3): 545–550. doi:10.1172/jci103832. PMC 293190 . PMID 13641405.

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