Solubility
59497
225405569
2008-07-13T14:42:09Z
Stan J Klimas
2562403
/* Applications */ - correct the link to indigo WP article that describes its solubility
'''Solubility''' is a characteristic [[physical property]] referring to the ability of a given substance, the [[solute]], to dissolve in a [[solvent]].<ref>''Atkins' Physical Chemistry'', 7th Ed. by Julio De Paula, P.W. Atkins ISBN 0198792859</ref> It is measured in terms of the maximum amount of solute [[solvation|dissolved]] in a solvent at [[chemical equilibrium|equilibrium]]. The resulting solution is called a saturated [[solution]]. Certain liquids are soluble in all proportions with a given solvent, such as [[ethanol]] in [[water]]. This property is known as [[miscible|miscibility]].<ref>Clugston M. and Fleming R. (2000), p.108</ref> Under certain conditions the [[solubility equilibrium|equilibrium solubility]] can be exceeded to give a so-called [[supersaturation|supersaturated]] solution, which is [[Metastability in molecules|metastable]].<ref>[http://cancerweb.ncl.ac.uk/cgi-bin/omd?metastable]: from Online Medical Dictionary, [[University of Newcastle Upon Tyne]]. </ref>
In a solution, the solvent is generally a liquid, which can be a pure substance or a [[mixture]].<ref>Yuen, C. (2003), ''Element, Compound and Mixture''</ref> The species that dissolves, the solute, can be a gas, another liquid, or a solid. Solubilities range widely, from infinitely soluble such as [[ethanol]] in [[water]], to poorly soluble, such as [[silver chloride]] in water. The term ''insoluble'' is often applied to poorly soluble compounds, although in some cases insolubility means that a compound is very poorly soluble.
==Molecular view==
Solubility occurs under dynamic equilibrium, which means that solubility results from the simultaneous and opposing processes of [[solvation|dissolution]] and [[Precipitation (chemistry)|precipitation]]. The solubility equilibrium occurs when the two processes proceed at the same rate.
The solubility equilibrium is relatively straightforward for [[covalent bond|covalent]] substances such as benzene. When dissolved in water, the benzene molecules remain intact but interact with and are generally surrounded by molecules of water. When, however, an [[ionic bond|ionic]] compound such as [[sodium chloride]] (NaCl) dissolves in water, the sodium chloride [[Crystal structure|lattice]] [[Dissociation (chemistry)|dissociates]] into individual ions that are [[solvation|solvated]] or surrounded by water molecules. Nonetheless, NaCl is said to dissolve in water, because evaporation of the solvent returns crystalline NaCl.
The term "dissolving" is sometimes applied to an irreversible [[chemical reaction]], as with iron in [[nitric acid]], but in such a case the thermodynamic concept of solubility does not apply.
When a solute dissolves, it may form several species in the solution. For example, an [[aqueous]] [[Suspension (chemistry)|suspension]] of [[Iron(II) hydroxide|ferrous hydroxide]], {{chem|Fe(OH)|2}}, will contain the series [{{chem|Fe(H|2|O)|6−x|(OH)|x}}]<sup>(2−x)+</sup> as well as other [[oligomer]]ic species. Furthermore, the solubility of ferrous hydroxide and the composition of its soluble components depends on [[pH]]. In general, solubility in the solvent phase can be given only for a specific solute which is thermodynamically stable, and the value of the solubility will include all the species in the solution (in the example above, all the iron-containing complexes).{{Fact|date=June 2008}}
==Factors affecting solubility==
Solubility is defined for specific [[phase (matter)|phases]]. For example, the solubility of [[aragonite]] and [[calcite]] in water are expected to differ, even though they are both [[polymorphs]] of [[calcium carbonate]], and have the same [[chemical formula]].
The solubility of one substance dissolving in another is determined by the balance of [[intermolecular force]]s between the solvent and solute, and the [[entropy]] change that accompanies the solvation. Factors such as temperature and pressure will alter this balance, thus changing the solubility.
Solubility may also strongly depend on the presence of other species dissolved in the solvent, for example, [[complex (chemistry)|complex-]]forming anions ([[ligand]]s) in liquids. Solubility will also depend on the excess or deficiency of a common ion in the solution, a phenomenon known as the [[common-ion effect]]. To a lesser extent, solubility will depend on the [[ionic strength]] of liquid solutions. The last two effects can be quantified using the equation for [[solubility equilibrium]].
Solubility (metastable) also depends on the physical size of the crystal or droplet of solute (or, strictly speaking, on the specific or molar surface area of the solute). For quantification, see the equation in the article on [[Solubility_equilibrium#Particle_size_effect|solubility equilibrium]]. For highly defective crystals, solubility may increase with the increasing degree of disorder. Both of these effects occur because of the dependence of solubility constant on the Gibbs energy of the crystal. The last two effects, although often difficult to measure, are of practical importance.{{Fact|date=July 2008}} For example, they provide the driving force for precipitate aging (the crystal size spontaneously increasing with time).
===Temperature===
[[Image:SolubilityVsTemperature.png|right|400px|Solubility of various salts as a function of temperature]]
[[Image:Temperature dependence solublity of solid in liquid water high temperature.svg|right|400px|Solubility of three salts in high-temperature liquid water]]
The solubility of a given solute in a given solvent typically depends on temperature. For around 95% of solids, the solubility increases with temperature from ambient to 100 °C.<ref name = hill>John W. Hill, Ralph H. Petrucci, ''General Chemistry'', 2nd edition, Prentice Hall, 1999.</ref> In liquid water at high temperatures, (e.g., that approaching the [[critical temperature]]), the solubility of ionic solutes tends to decrease due to the change of properties and structure of liquid water; the lower [[dielectric constant]] results in a less [[polar solvent]].
[[Gas]]eous solutes exhibit more complex behavior with temperature. As the temperature is raised, gases usually become less soluble in water, but more soluble in organic solvents.<ref name=hill/>
The chart shows solubility curves for some typical solid inorganic [[salt]]s.<ref> Data taken from the ''Handbook of Chemistry and Physics'', 27th edition, Chemical Rubber Publishing Co., Cleveland, Ohio, 1943.</ref> Many salts behave like [[barium nitrate]] and disodium hydrogen arsenate, and show a large increase in solubility with temperature. Some solutes (e.g. NaCl in water) exhibit solubility which is fairly independent of temperature. A few, such as cerium(III) sulfate, become less soluble in water as temperature increases. This is sometimes referred to as "retrograde" or "inverse" solubility. Occasionally, a more complex pattern is observed, as with [[sodium sulfate]], where the less soluble deca[[hydrate]] crystal loses [[water of crystallization]] at 32 °C to form a more soluble [[anhydrous]] phase.{{Fact|date=July 2008}}
The solubility of [[organic compounds]] nearly always increases with temperature. The technique of [[recrystallization]], used for purification of solids, depends on a solute's different solubilities in hot and cold solvent. A few exceptions exist, such as certain [[cyclodextrin]]s.<ref>{{cite journal|title = A highly water-soluble 2+1 b-cyclodextrin–fullerene conjugate | author = Salvatore Filippone, Frank Heimanna and André Rassat | journal = [[Chem. Commun.]] | volume = 2002 | pages = 1508–1509 | doi = 10.1039/b202410a | year = 2002}}</ref>
===Pressure===
For condensed phases (solids and liquids), the pressure dependence of solubility is typically weak and usually neglected in practice. Assuming an ideal solution, the dependence can be quantified as:
:<math> \left(\frac{\partial \ln N_i}{\partial P} \right)_T = -\frac{V_{i,aq}-V_{i,cr}} {RT} </math>
where the index i iterates the components, N<sub>i</sub> is the mole fraction of the i<sup>th</sup> component in the solution, P is the pressure, the index T refers to constant temperature, V<sub>i,aq</sub> is the [[partial molar volume]] of the i<sup>th</sup> component in the solution, V<sub>i,cr</sub> is the partial molar volume of the i<sup>th</sup> component in the dissolving solid, and R is the [[universal gas constant]]<ref>E.M.Gutman, "Mechanochemistry of Solid Surfaces", World Scientific Publishing Co., 1994.</ref>.
==Solubility of gases==
[[Henry's law]] is used to quantify the solubility of gases in solvents. The solubility of a gas in a solvent is directly proportional to the [[partial pressure]] of that gas above the solvent. This relationship is written as:
:<math> p = kc \,</math>
where k is a temperature-dependent constant (for example, 769.2 [[litre|L]]•[[Atmosphere (unit)|atm]]/[[Mole (unit)|mol]] for [[dioxygen]] (O<sub>2</sub>) in water at 298 K), p is the partial pressure (atm), and c is the [[concentration]] of the dissolved gas in the liquid (mol/L).
==Polarity==
A popular [[aphorism]] used for predicting solubility is "''like dissolves like''".<ref>Kenneth J. Williamson, ''Macroscale and Microscale Organic Experiments'', p40, 2nd edition, D. C, Heath, Lexington, Mass., 1994.</ref> This indicates that a solute will dissolve best in a solvent that has a similar [[Chemical polarity|polarity]] to itself. This is a rather simplistic view, since it ignores many solvent-solute interactions, but it is a useful rule-of-thumb. For example, a very polar ([[hydrophile|hydrophilic]]) solute such as [[urea]] is very soluble in highly polar water, less soluble in fairly polar [[methanol]], and practically insoluble in non-polar solvents such as [[benzene]]. In contrast, a non-polar or [[lipophilicity|lipophilic]] solute such as [[naphthalene]] is insoluble in water, fairly soluble in methanol, and highly soluble in non-polar benzene.<ref>Data taken from the ''Merck Index'', 7th edition, Merck & Co., 1960.</ref>
Liquid solubilities also generally follow this rule. Lipophilic plant oils, such as olive oil and palm oil, dissolve in non-polar solvents such as alkanes, but are less soluble in polar liquids such as water.
Synthetic chemists often exploit differences in solubilities to separate and purify compounds from reaction mixtures, using the technique of [[liquid-liquid extraction]].
==Rate of dissolution==
[[Solvation|Dissolution]] is not always an instantaneous process. It is fast when salt and sugar dissolve in water but much slower for a tablet of [[aspirin]] or a large crystal of hydrated [[copper(II) sulfate]]. The speed at which a solid dissolves may depend on its crystalline properties (i.e. whether it is crystalline or[[amorphous]]), and the crystal size) and the presence of [[Polymorphism (materials science)|polymorphism]]. This is important in many practical systems, for example in designing methods for controlled [[drug delivery]]. Critically, the dissolution rate depends on the presence of mixing and other factors that determine the degree of undersaturation in the liquid solvent film immediately adjacent to the solid solute crystal. In some cases, solubility equilibria can take a long time to establish (hours, days, months, or many years; depending on the nature of the solute and other factors). In practice, it means that the amount of solute in a solution is not always determined by its thermodynamic solubility, but may depend on kinetics of dissolution (or precipitation).
The rate of dissolution and solubility should not be confused as they are different concepts, kinetic and thermodynamic, respectively.
==Quantification of solubility==
Solubility is commonly expressed as a concentration, either by mass (g of solute per kg of solvent, g per dL (100 mL) of solvent), [[molarity]], [[molality]], mole fraction or other similar descriptions of concentration. The maximum equilibrium amount of solute that can dissolve per amount of solvent is the solubility of that solute in that solvent under the specified conditions. The advantage of expressing solubility in this manner is its simplicity, while the disadvantage is that it can strongly depend on the presence of other species in the solvent (for example, the common ion effect).
[[Solubility constant]]s are used to describe saturated solutions of ionic compounds of relatively low solubility (see [[solubility equilibrium]]). The solubility constant is a special case of an [[equilibrium constant]]. It describes the balance between dissolved ions from the salt and undissolved salt. The solubility constant is also "applicable" (i.e. useful) to [[precipitation (chemistry)|precipitation]], the reverse of the dissolving reaction. As with other equilibrium constants, [[temperature]] can affect the numerical value of solubility constant. The solubility constant is not as simple as solubility, however the value of this constant is generally independent of the presence of other species in the solvent.
The [[Flory-Huggins solution theory]] is a theoretical model describing the solubility of polymers. The [[Hansen Solubility Parameters]] and the [[Hildebrand solubility parameter]]s are empirical methods for the prediction of solubility. It is also possible to predict solubility from other physical constants such as the [[enthalpy of fusion]].
The [[partition coefficient]] ([[Log P]]) is a measure of differential solubility of a compound in a [[hydrophobe|hydrophobic]] solvent ([[octanol]]) and a [[hydrophile|hydrophilic]] solvent ([[water]]). The logarithm of these two values enables compounds to be ranked in terms of hydrophilicity (or hydrophobicity).
==Applications==
Solubility is of fundamental importance in a large number of scientific disciplines and practical applications, ranging from ore processing, to the use of medicines, and the transport of pollutants.
Solubility is often said to be one of the "characteristic properties of a substance," which means that solubility is commonly used to describe the substance, to indicate a substance's polarity, to help to distinguish it from other substances, and as a guide to applications of the substance. For example, [[Indigo_dye#Chemical_properties|indigo]] is described as "insoluble in water, alcohol, or ether but soluble in chloroform, nitrobenzene, or concentrated sulfuric acid".{{Fact|date=July 2008}}
Solubility of a substance is useful when separating mixtures. For example, a mixture of salt ([[sodium chloride]]) and silica may be separated by dissolving the salt in water, and filtering off the undissolved silica. The synthesis of chemical compounds, by the milligram in a laboratory, or by the ton in industry, both make use of the relative solubilities of the desired product, as well as unreacted starting materials, byproducts, and side products to achieve separation.
Another example of this is the synthesis of [[benzoic acid]] from [[phenylmagnesium bromide]] and [[dry ice]]. Benzoic acid is more soluble in an organic solvent such as [[dichloromethane]] or [[diethyl ether]], and when shaken with this organic solvent in a [[separatory funnel]], will preferentially dissolve in the organic layer. The other reaction products, including the magnesium bromide, will remain in the aqueous layer, clearly showing that separation based on solubility is achieved. This process, known as [[liquid-liquid extraction]], is an important technique in [[synthetic chemistry]].
==Solubility of ionic compounds in water ==
Some ionic compounds ([[salts]]) dissolve in water, which arises because of the attraction between positive and negative charges (see: [[solvation]]). For example, the salt's positive ions (i.e. Ag<sup>+</sup>) attract the partially-negative oxygens in H<sub>2</sub>O. Likewise, the salt's negative ions (i.e. Cl<sup>−</sup>) attract the partially-positive hydrogens in in H<sub>2</sub>O. Note: oxygen is partially-negative because it is more [[electronegative]] than hydrogen, and vice-versa (see: [[chemical polarity]]).
:AgCl<sub>(s)</sub> <math>\overrightarrow{\leftarrow}</math> Ag<sup>+</sup><sub>(aq)</sub> + Cl<sup>−</sup><sub>(aq)</sub>
However, there is a limit to how much salt can be dissolved in a given volume of water. This amount is given by the [[solubility product]], K<sub>sp</sub>. This value depends on the type of salt (AgCl vs. NaI, for example), temperature, and the [[common ion effect]].
One can calculate the amount of AgCl that will dissolve in 1 liter of water, some algebra is required.
:K<sub>sp</sub> = [Ag<sup>+</sup>] × [Cl<sup>−</sup>] (definition of solubility product)
:K<sub>sp</sub> = 1.8 × 10<sup>−10</sup> (from a table of solubility products)
[Ag<sup>+</sup>] = [Cl<sup>−</sup>], in the absence of other silver or chloride salts,
:[Ag<sup>+</sup>]<sup>2</sup> = 1.8 × 10<sup>−10</sup>
:[Ag<sup>+</sup>] = 1.34 × 10<sup>−5</sup>
The result: 1 liter of water can dissolve 1.34 × 10<sup>−5</sup> [[mole (unit)|moles]] of AgCl<sub>(s)</sub> at room temperature. Compared with other types of salts, AgCl is poorly soluble in water. In contrast, table salt (NaCl) has a higher K<sub>sp</sub> and is, therefore, more soluble.
{{main|Solubility chart}}
<center><table class="wikitable">
<tr><th>Soluble</th><th>Insoluble</th></tr>
<tr><td>[[Alkali metal|Group I]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds</td><td>[[carbonate]]s (except [[Alkali metal|Group I]], [[Ammonium|NH<sub>4</sub><sup>+</sup>]] and [[uranyl]] compounds)</td></tr>
<tr><td>[[nitrate]]s</td><td>[[sulfite]]s (except [[Alkali metal|Group I]] and [[Ammonium|NH<sub>4</sub><sup>+</sup> compounds]])</td></tr>
<tr><td>[[acetate]]s (ethanoates)</td><td>[[phosphate]]s (except [[Alkali metal|Group I]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds)</td></tr>
<tr><td>[[chloride]]s, [[bromide]]s and [[iodide]]s (except [[Silver|Ag<sup>+</sup>]], [[Lead|Pb<sup>2+</sup>]], [[Copper|Cu<sup>+</sup>]] and [[Mercury (element)|Hg<sub>2</sub><sup>2+</sup>]])</td><td>[[hydroxide]]s and [[oxide]]s (except [[Alkali metal|Group I]], [[Ammonium|NH<sub>4</sub><sup>+</sup>]], [[Barium|Ba<sup>2+</sup>]], [[Strontium|Sr<sup>2+</sup>]] and [[Thallium|Tl<sup>+</sup>]])</td></tr>
<tr><td>[[sulfate]]s (except [[Silver|Ag<sup>+</sup>]], [[Lead|Pb<sup>2+</sup>]], [[Barium|Ba<sup>2+</sup>]], [[Strontium|Sr<sup>2+</sup>]] and [[Calcium|Ca<sup>2+</sup>]])</td><td>[[sulfide]]s (except [[Alkali metal|Group I]], [[Alkaline earth metal|Group II]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds)</td></tr>
</table>
</center>
==Solubility of organic compounds==
The principle outlined above under [[#Polarity|polarity]], that ''like dissolves like,'' is the usual guide to solubility with organic systems. For example, [[petroleum jelly]] will dissolve in [[gasoline]]; both of which are lipophilic. This is because petroleum jelly consists of long carbon chains, as does the gasoline. It will not, on the other hand, dissolve in alcohol or water, since the polarity of these solvents is too high. Sugar will not dissolve in gasoline, since sugar is too polar in comparison with gasoline. A mixture of gasoline and sugar can therefore be separated by [[filtration]], or [[solvent extraction|extraction]] with water.
==Solid solution==
This term is often used in the field of [[metallurgy]] to refer to the extent that an [[alloy]]ing element will dissolve into the [[base metal]] without forming a separate phase. The solubility line (or curve) is the line (or lines) on a [[phase diagram]] which give the limits of solute addition. That is, the lines show the maximum amount of a component that can be added to another component and still be in [[solid solution]]. In microelectronic fabrication, solid solubility refers to the maximum concentration of impurities one can place into the substrate.
==Incongruent dissolution==
Many substances dissolve congruently, i.e., the composition of the solid and the dissolved solute stoichiometrically match. However, some substances may dissolve [[Incongruent transition|incongruently]], whereby the composition of the solute in solution does not match that of the solid. This is accompanied by alteration of the "primary solid" and possibly formation of a secondary solid phase. However, generally, some primary solid also remains and a complex solubility equilibrium establishes. For example, dissolution of [[albite]] may result in formation of [[gibbsite]].<ref>O.M.Saether & P. de Caritat (ed.) "Geochemical processes, weathering and groundwater recharge in catchments", Taylor & Francis, Rotterdam, 1997, page 6.</ref>
:NaAlSi<sub>3</sub>O<sub>8</sub>(s) + H<sup>+</sup> + 7H<sub>2</sub>O = Na<sup>+</sup> + Al(OH)<sub>3</sub>(s) + 3H<sub>4</sub>SiO<sub>4</sub>.
In this case, the solubility of albite is expected to depend on the solid-to-solvent ratio. This kind of solubility is of great importance in geology, where it results in formation of [[metamorphic rock]]s.
== See also ==
{{wiktionarypar|soluble|solubility}}
* [[Solubility equilibrium]]
* [[Biopharmaceutics Classification System]]
* [[Hot water extraction]]
== External links ==
*[http://www.vcclab.org/lab/alogps ALOGPS] interactive calculation of aqueous solubility of compounds at [http://www.vcclab.org Virtual Computational Chemistry Laboratory] using several algorithms
*[http://q-lead.com/cnt/LogS/ QUANTUM] web based calculation of aqueous and DMSO solubility of compounds QUANTUM web based prediction of aqueous and DMSO solubility of compounds
*[http://www.acdlabs.com/products/phys_chem_lab/aqsol/ ACD/Solubility DB] aqueous solubility prediction
==References==
{{reflist}}
{{Chemical solutions}}
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