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Osmotic concentration
Osmotic concentration
Osmotic concentration
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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 particles on dissociation of osmotically active material (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.

An ORS sachet with the osmolarity of its components

Unit

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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 one thousandth of an osmole. A microosmole (μOsm) (also spelled micro-osmole) is one millionth of an osmole.

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 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] [Note: NaCl does not dissociate completely in water at standard temperature and pressure, so the solution will be composed of Na+ ions, Cl- ions, and some NaCl molecules, with actual osmolality = Na+ concentration x 1.75]

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

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The osmolarity of a solution, given in osmoles per liter (osmol/L) is 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.

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

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Osmolarity and tonicity are related but distinct concepts. Thus, the terms ending in -osmotic (isosmotic, hyperosmotic, hypoosmotic) 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.

In medicine

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Plasma osmolarity vs. osmolality

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

Osmolarity = osmolality × (ρsol − ca)

where:

  • ρsol is the density of the solution in g/ml, which is 1.025 g/ml for blood plasma.[5]
  • ca 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).

where mi is the molality of component i.

Plasma osmolarity/osmolality is important for keeping proper electrolytic balance in the blood stream. Normally, sodium is the major contributor to plasma (or serum) osmolality; glucose and urea also contribute.[6] 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.[7]

Hyperosmolarity and hypoosmolarity

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A concentration of an osmotically active substance is said to be hyperosmolar if a high concentration causes a change in osmotic pressure in a tissue, organ, or system. Hypernatremia, diabetic ketoacidosis, and uremia are pathologic causes of hyperosmolarity.[6] Similarly, it is said to be hypoossmolar if the osmolarity, or osmatic concentration, is too low. For example, if the osmolarity of parenteral nutrition is too high, it can cause severe tissue damage.[8] One example of a condition associated with hypoosmolarity is water intoxication.[9]

See also

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References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Osmotic concentration

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Osmotic concentration refers to the total number of solute particles dissolved in a solution, independent of their chemical identity, and is a key determinant of osmotic pressure and water movement across semipermeable membranes. It is quantified either as osmolarity, the number of osmoles per liter of solution, or osmolality, the number of osmoles per kilogram of solvent, with both terms often used somewhat interchangeably in biological contexts despite their precise differences. This concentration drives osmosis, the passive diffusion of water from a region of lower solute concentration (higher water potential) to higher solute concentration (lower water potential), which is essential for maintaining cellular integrity and fluid balance in living organisms. The distinction between osmolarity and osmolality arises from their reference frames: osmolarity accounts for the volume of the entire solution, which can vary with due to solute-solvent interactions, whereas osmolality is based on the of the alone and remains more stable across changes. In practice, osmolality is preferred in clinical and physiological measurements because it provides a temperature-independent assessment of solute concentration. For instance, a 1 molal solution of a non-dissociating solute like glucose yields 1 osmol/kg, but a dissociating solute like (NaCl) produces approximately 2 osmoles/kg due to its ionization into Na⁺ and Cl⁻ ions. Osmotic concentration is typically expressed in osmoles (Osm) or milliosmoles (mOsm), with normal human serum osmolality ranging from 275 to 295 mOsm/kg, reflecting a balance of electrolytes, , and glucose. It is measured indirectly through , such as or reduction, using an that detects minute temperature changes in a small sample (e.g., 10 µL at 37°C). The (π) generated by a solute can be calculated using the : π = iCRT, where i is the (number of particles per molecule), C is the , R is the , and T is the absolute temperature. In biology, osmotic concentration is critical for processes like nutrient absorption in the intestines, kidney function in regulating body fluids, and cellular responses to environmental changes. For example, in hypertonic solutions (higher osmotic concentration outside the cell), water exits the cell, causing shrinkage (crenation in red blood cells); in hypotonic solutions (lower external concentration), water enters, leading to swelling or lysis; while isotonic conditions maintain equilibrium. Clinically, deviations in osmotic concentration can indicate disorders such as dehydration, syndrome of inappropriate antidiuretic hormone secretion (SIADH), or diabetes insipidus, guiding diagnostic and therapeutic interventions.

Basic Concepts

Definition

Osmotic concentration refers to the measure of the total number of solute particles in a solution, quantified either as the number of osmoles per liter of solution (osmolarity, osmol/L) or per of solvent (osmolality, osmol/kg), which determines the solution's capacity to exert osmotic effects across a . An osmole represents the amount of solute that contributes one mole of particles to the osmotic activity, accounting for dissociation or association behaviors that affect the effective particle count. This concentration drives , the passive movement of across a from a region of lower osmotic concentration (higher ) to one of higher osmotic concentration (lower ), aiming to equalize the solute particle distribution on both sides. The resulting osmotic influences cellular and physiological processes by regulating , with higher osmotic concentrations creating stronger pulls for water influx. The concept originated in the late through the work of Dutch chemist , who in demonstrated that in dilute solutions follows principles analogous to laws, establishing osmotic concentration as a key colligative property dependent on solute particle number rather than identity. 's foundational contributions, recognized by the 1901 , linked osmotic behavior to thermodynamic equilibria in solutions. For illustration, pure has an osmotic concentration of 0 osmol/L, exerting no osmotic pull, whereas a 0.9% (NaCl) solution, commonly used as normal saline, has an approximate osmotic concentration of 0.3 osmol/L due to the dissociation of NaCl into two ions per molecule. This difference highlights how even modest solute additions can generate significant osmotic gradients.

Units of Measurement

The osmole (osmol), a non-SI unit, represents the amount of solute that contributes one mole of osmotically active particles to a solution, accounting for dissociation or association of the solute molecules. Osmotic concentration is quantified using this unit, expressed either as osmolality (osmoles per of ) or osmolarity (osmoles per liter of solution), which reflect the total particle concentration influencing osmotic behavior. In biological and medical contexts, subunits such as the milliosmole (mosmol) are commonly employed due to the relatively low concentrations encountered, with normal human plasma osmolality ranging from 275 to 295 mosmol/kg. This scale allows precise measurement of physiological fluids, where even small deviations can indicate imbalances in solute distribution. To relate osmotic concentration to standard molarity (moles per liter), the van't Hoff factor (i) is used, which represents the average number of particles per solute molecule; osmolarity is thus approximated as i multiplied by the molarity for ideal solutions, without deriving the full osmotic pressure equation. Although the mole is the for , osmotic concentrations are not formally SI-derived but align with mmol/kg for osmolality in clinical reporting; osmolality is preferred over osmolarity in precise biological measurements because it is mass-based, remaining unaffected by or variations that influence solution volume.

Solution Properties

Types of Solutes

Solutes contributing to osmotic concentration are broadly classified into non-electrolytes and electrolytes based on their dissociation behavior in solution, which determines the van't Hoff factor (i) used in calculating effective particle concentration. Non-electrolytes, such as glucose, do not dissociate into ions upon dissolution and thus have a van't Hoff factor of i=1, meaning a 1 mol/L solution of glucose yields an osmotic concentration of 1 osmol/L. These solutes contribute directly to the total osmole count without increasing the number of particles beyond the original molecules. In contrast, electrolytes dissociate into multiple ions, resulting in a van't Hoff factor greater than 1; for example, (NaCl) ideally dissociates into Na⁺ and Cl⁻ ions, giving i=2 for a 1 mol/L solution and thus 2 osmol/L. Common physiological electrolytes include ions like Na⁺, Cl⁻, and K⁺, which are prevalent in biological fluids and significantly influence osmotic concentration due to their dissociation. Solutes are further categorized by membrane permeability, which affects their effective contribution to osmotic gradients across semi-permeable barriers: impermeant solutes cannot readily cross the , maintaining a persistent concentration difference, while permeant solutes can diffuse across, potentially reducing their net osmotic effect over time. In biological solutions, serves as a permeant non-electrolyte (i=1) that equilibrates across cell membranes, whereas acts as an impermeant non-electrolyte (i=1) commonly used in osmotic therapies due to its restricted passage through certain barriers. These classifications underpin the expression of osmotic concentration in units like osmol/L, reflecting the total effective solute particles.

Osmolality versus Osmolarity

Osmolality refers to the concentration of osmoles of solute per of , providing a measure that is independent of the solution's volume and more suitable for precise assessments in varying conditions. In contrast, osmolarity denotes the number of osmoles of solute per liter of solution, which relies on the volume of the entire mixture and is thus influenced by factors such as and solute addition. The primary difference between osmolality and osmolarity lies in their basis of measurement: osmolality uses the mass of the , making it less affected by expansions or contractions caused by solutes or changes, which is particularly advantageous for biological fluids where is not constant. Osmolarity, assuming a uniform , simplifies theoretical calculations but introduces inaccuracies in non-ideal solutions, as the addition of solutes can alter the total without proportionally changing the solvent mass. This makes osmolality the preferred metric in physiological contexts, where precise solute-solvent interactions are critical. For dilute aqueous solutions, osmolality and osmolarity are approximately equal due to the near-unity density of water (about 0.997 g/mL at 25°C), with the difference typically negligible (less than 1%). In practice, osmolality is directly measured using freezing point depression osmometry, which detects the colligative property lowering of the solvent's freezing point proportional to solute particles. Osmolarity, however, is typically estimated through calculations from known solute concentrations, such as electrolytes in serum panels, offering a convenient approximation when direct measurement is unnecessary.

Osmotic Phenomena

Tonicity

Tonicity refers to the ability of an extracellular solution to influence the volume of a cell by causing net movement of water across the cell membrane through osmosis, determined primarily by the concentration of solutes that cannot permeate the membrane. Unlike total osmotic concentration, tonicity focuses on the effective osmotic gradient created by impermeant solutes, as permeant solutes equilibrate across the membrane without sustaining a lasting water flux. The key determinant of tonicity is the permeability of the to the solutes in the solution; only impermeant solutes, such as Na⁺ and Cl⁻ in NaCl, generate an osmotic gradient that affects cell volume, whereas permeant solutes like do not contribute significantly because they diffuse freely into the cell. For instance, a solution of 300 mM is isosmotic to intracellular fluid but hypotonic because permeates the , leading to transient cell swelling without long-term change. In contrast, 150 mM NaCl creates an isotonic environment relative to typical intracellular conditions, resulting in no net water movement. Solutions are classified by as isotonic, hypertonic, or hypotonic based on their effect on cell volume compared to the intracellular environment. Isotonic solutions, such as 0.9% NaCl (normal saline), result in no net movement and maintain cell volume, making them suitable for intravenous administration to avoid cellular disruption. Hypertonic solutions have a higher concentration of impermeant solutes than the cell interior, causing to exit the cell and leading to shrinkage ( in blood cells). Hypotonic solutions, with lower impermeant solute concentrations, drive into the cell, resulting in swelling and potential if unchecked. Examples of isotonic intravenous fluids include lactated , which approximates the composition of plasma and prevents net water shifts in human cells during fluid therapy. These classifications are critical for understanding how solutions interact with biological membranes, emphasizing the role of solute permeability over total osmolality.

Osmotic Pressure

Osmotic pressure, denoted as π, is defined as the minimum pressure that must be applied to a solution to prevent the inward flow of pure solvent across a , thereby halting . This colligative property arises directly from the osmotic concentration of the solute and is independent of the solute's chemical identity, depending solely on the number of solute particles per unit volume. The quantitative relationship between osmotic pressure and osmotic concentration is given by the , derived in 1887: π=iCRT\pi = iCRT where ii is the representing the number of particles into which a solute dissociates (e.g., i=2i = 2 for NaCl), CC is the osmotic concentration in osmol/L, RR is the universal (8.314 J/mol·K), and TT is the absolute temperature in . This equation stems from an to the PV=nRTPV = nRT: van 't Hoff conceptualized the solute particles in solution as exerting a on the similar to gas molecules confined in a volume, where the effective "volume" is 1 L and the number of moles nn corresponds to CC. At equilibrium, the applied pressure balances the tendency of solvent to flow due to the chemical potential difference, leading to the linear proportionality π∝C\pi \propto C for dilute, ideal solutions. In non-ideal solutions, particularly at higher concentrations, deviations from the occur due to intermolecular interactions between solute particles, which reduce the effective activity of the solutes. These deviations are accounted for by incorporating activity coefficients (γ\gamma), modifying the equation to π=iγCRT\pi = i \gamma C R T, where γ<1\gamma < 1 reflects the non-ideality; for electrolytes, a mean ionic activity coefficient is used. Such corrections are essential in concentrated systems, as solute-solute attractions or repulsions alter the colligative behavior predicted by the ideal model. Osmotic pressure is measured using osmometry techniques, which exploit colligative properties. Membrane osmometry directly applies hydrostatic pressure across a semipermeable membrane to achieve equilibrium and quantify π, suitable for high-molecular-weight solutes like polymers. Vapor pressure osmometry, an indirect method, measures the temperature difference caused by solvent evaporation onto a thermocouple in the presence of the solution, correlating it to osmotic concentration via the colligative depression of vapor pressure. These methods provide accurate assessments, with membrane techniques preferred for absolute pressure values in dilute solutions up to several atmospheres.

Biological and Medical Applications

Physiological Role

Osmotic concentration plays a critical role in regulating cell volume at the cellular level by maintaining osmotic balance between the cytoplasm and the extracellular environment. In mammalian cells, the cytoplasmic osmotic concentration is typically around 300 mOsm/kg, which helps prevent excessive water influx or efflux across semi-permeable membranes, thereby preserving cellular integrity and function. This balance is essential for processes such as nutrient uptake and waste removal, where deviations can lead to swelling or shrinkage, but under normal conditions, it ensures stable volume through mechanisms like tonicity control. At the organismal level, osmotic concentration is vital for kidney function, particularly in the countercurrent multiplier system of the loop of Henle, which establishes a medullary osmotic gradient to concentrate urine and conserve water. This system allows the kidney to produce urine with osmolalities up to 1200 mOsm/kg, enabling efficient reabsorption of water from the filtrate while excreting waste solutes. By modulating osmotic gradients in the renal medulla, this mechanism maintains overall fluid and electrolyte homeostasis in vertebrates. In plants, osmotic concentration in central vacuoles drives turgor pressure, which provides structural support and facilitates cell expansion for growth. Solutes accumulated in the vacuole create an osmotic gradient that draws water inward, generating hydrostatic pressure against the cell wall and maintaining plant rigidity. Similarly, bacteria employ osmoregulation to counteract environmental osmotic shifts by synthesizing or transporting compatible solutes, such as glycine betaine or trehalose, to stabilize cytoplasmic osmotic concentration and preserve turgor for survival in fluctuating salinities. These osmotic mechanisms exhibit evolutionary conservation from prokaryotes to eukaryotes, reflecting adaptations for survival in diverse osmotic environments across the tree of life. Fundamental strategies, such as solute accumulation and membrane permeability regulation, have been retained and refined over billions of years, underscoring their foundational role in cellular homeostasis.

Clinical Measurement and Disorders

In clinical practice, plasma osmolality is measured using direct methods such as freezing-point depression osmometry (cryoscopy), which determines the colligative properties of solutes by measuring the temperature at which the solution freezes, or vapor pressure osmometry, which assesses the vapor pressure lowering caused by non-volatile solutes. These direct techniques provide accurate total osmolality without relying on specific solute concentrations. Indirect methods, commonly used in laboratories, calculate osmolality from serum electrolytes, glucose, and urea concentrations via formulas such as 2[Na⁺] + [glucose]/18 + [BUN]/2.8 (where concentrations are in mg/dL for glucose and BUN), offering a cost-effective approximation but potentially underestimating unmeasured osmoles. The normal range for plasma osmolality in adults is 275–295 mosmol/kg, reflecting a tightly regulated balance essential for cellular function. An osmolal gap, calculated as the difference between directly measured and indirectly calculated osmolality, normally falls below 10 mosmol/kg; elevations beyond this threshold indicate unmeasured osmoles, such as those from toxins like methanol or ethylene glycol, aiding in the diagnosis of intoxications. Hyponatremia, characterized by hypoosmolarity (plasma osmolality <275 mosmol/kg), often arises from causes like the syndrome of inappropriate antidiuretic hormone secretion (SIADH), where excessive antidiuretic hormone leads to water retention and dilutional sodium loss. Symptoms can include severe neurological effects such as seizures, particularly when sodium levels drop rapidly below 120 mEq/L, necessitating prompt intervention to prevent cerebral edema. In contrast, hyperosmolar hyperglycemic state (HHS), a complication of type 2 diabetes, involves marked hyperosmolarity (often >320 mosmol/kg) due to severe and , without significant . Treatment prioritizes aggressive with intravenous isotonic saline to correct volume depletion and gradually lower osmolality, alongside insulin therapy to resolve . Post-2020 research has highlighted osmolality's role in severe cases, where and cytokine storms contribute to dysregulated sodium and levels, elevating and correlating with higher mortality risk in over 1,300 analyzed patients. Updated guidelines from the European Society of Endocrinology emphasize the dangers of rapid correction, recommending limits of 10 mmol/L in 24 hours and 18 mmol/L in 48 hours to avoid osmotic demyelination syndrome, particularly in chronic cases.

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