

The solvent is generally a liquid, which can be a pure substance or a mixture.[1] One also speaks of solid solution, but rarely of solution in a gas (see vapor-liquid equilibrium instead).

The extent of solubility ranges widely, from infinitely soluble (fully miscible[2] ) such as ethanol in water, to poorly soluble, such as silver chloride in water. The term insoluble is often applied to poorly or very poorly soluble compounds.


Solubility is not to be confused with the ability to dissolve or liquefy a substance, because they might occur not only because of dissolution but also because of a chemical reaction. For example, zinc is insoluble in hydrochloric acid, but does dissolve in it by chemical reaction into zinc chloride and hydrogen, where zinc chloride is then soluble in hydrochloric acid.


Molecular view


The term solubility is also used in some fields where the solute is altered by solvolysis. For example, many metals and their oxides are said to be "soluble in hydrochloric acid," whereas the aqueous acid degrades the solid to irreversibly give soluble products. It is also true that most ionic solids are degraded by polar solvents, but such processes are reversible. In those cases where the solute is not recovered upon evaporation of the solvent the process is referred to as solvolysis. The thermodynamic concept of solubility does not apply straightforwardly to solvolysis.

When a solute dissolves, it may form several species in the solution. For example, an aqueous suspension of ferrous hydroxide, Fe(OH)2, will contain the series [Fe(H2O)6 − x(OH)x](2 − x)+ as well as other oligomeric 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).

Factors affecting solubility




For a solid that dissolves in a redox reaction, solubility is expected to depend on the potential (within the range of potentials under which the solid remains the thermodynamically stable phase). For example, solubility of gold in high-temperature water is observed to be almost an order of magnitude higher when the redox potential is controlled using a highly-oxidizing Fe3O4-Fe2O3 redox buffer than with a moderately-oxidizing Ni-NiO buffer.[4]


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. 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. For example, they provide the driving force for precipitate aging (the crystal size spontaneously increasing with time).


The solubility of a given solute in a given solvent typically depends on temperature. For many solids dissolved in liquid water, the solubility increases with temperature up to 100 °C.[5] 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.

Gaseous solutes exhibit more complex behavior with temperature. As the temperature is raised, gases usually become less soluble in water (to minimum which is below 120 °C for most permanent gases[6]), but more soluble in organic solvents.[5]

The chart shows solubility curves for some typical solid inorganic salts (temperature is in degrees Celsius).[7] 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 temperature dependence is sometimes referred to as "retrograde" or "inverse" solubility. Occasionally, a more complex pattern is observed, as with sodium sulfate, where the less soluble decahydrate crystal loses water of crystallization at 32 °C to form a more soluble anhydrous phase.

Temperature dependence solublity of solid in liquid water high temperature.svg







Solubility of gases


 p = k_{\rm H}\, c


In the presence of small bubbles, the solubility of the gas does not depend on the bubble radius in any other way than through the effect of the radius on pressure (see Young–Laplace equation)[10]. This is unlike small crystals, for which the solubility increases.


A popular aphorism used for predicting solubility is "like dissolves like".[11] This statement indicates that a solute will dissolve best in a solvent that has a similar polarity to itself. This view is rather simplistic, since it ignores many solvent-solute interactions, but it is a useful rule of thumb. For example, a very polar (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 lipophilic solute such as naphthalene is insoluble in water, fairly soluble in methanol, and highly soluble in non-polar benzene.[12]

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.


Insolubility and spontaneous phase separation does not mean that dissolution is disfavored by enthalpy. Quite the contrary, in the case of water and hydrophobic substances, hydrophobic hydration is reasonably exothermic and enthalpy alone should favor it. It appears that entropic factors — the reduced freedom of movement of water molecules around hydrophobic molecules — lead to an overall hydrophobic effect.

Rate of dissolution


The rate of dissolution and solubility should not be confused as they are different concepts, kinetic and thermodynamic, respectively. The solubilization kinetics, as well as apparent solubility can be improved after complexation of an active ingredient with cyclodextrin. This can be used in the case of drug with poor solubility.[13]

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 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 is described as "insoluble in water, alcohol, or ether but soluble in chloroform, nitrobenzene, or concentrated sulfuric acid".



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 (e.g. Ag+) attract the partially-negative oxygens in H2O. Likewise, the salt's negative ions (e.g. Cl) attract the partially-positive hydrogens in H2O. Note: oxygen is partially-negative because it is more electronegative than hydrogen, and vice-versa (see: chemical polarity).

AgCl(s) is in equilibrium with Ag+(aq) + Cl(aq)





[Ag+]2 = 1.8 × 10−10
[Ag+] = 1.34 × 10−5


Soluble Insoluble
Group I and NH4+ compounds Carbonates (Except Group I, NH4+ and uranyl compounds)
Nitrates Sulfites (Except Group I and NH4+ compounds)
Acetates (Ethanoates) (Except Ag+ compounds) Phosphates (Except Group I and NH4+ compounds)

Solubility of organic compounds


Solubility in non-aqueous solvents

Most publicly available solubility values are those for solubility in water.[14] The reference also lists some for non-aqueous solvents. Solubility data for non-aqueous solvents is currently being collected via an open notebook science crowdsourcing project.[15][16]

Solid solution

This term is often used in the field of metallurgy to refer to the extent that an alloying 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 the solid's crystalline structure, the 'solute' element can either take the place of the matrix within the lattice (a substitutional position, for example: chromium in iron) or can take a place in a space between the lattice points (an interstitial position, for example: carbon in iron).


Incongruent dissolution




See also


  1. Yuen, C. (2003), Element, Compound and Mixture
  2. Clugston M. and Fleming R. (2000), p.108
  3. from Online Medical Dictionary, University of Newcastle Upon Tyne.
  4. I.Y. Nekrasov, "Geochemistry, Mineralogy and Genesis of Gold Deposits", Taylor & Francis, 1996, p.135-136
  5. 5.0 5.1 John W. Hill, Ralph H. Petrucci, General Chemistry, 2nd edition, Prentice Hall, 1999.
  6. P. Cohen (editor), "The ASME handbook on Water Technology for Thermal Power Systems", The American Society of Mechanical Engineers, 1989, page 442.
  8. Salvatore Filippone, Frank Heimanna and André Rassat (2002). "A highly water-soluble 2+1 b-cyclodextrin–fullerene conjugate". Chem. Commun. 2002: 1508–1509. doi:10.1039/b202410a. 
  9. E.M.Gutman, "Mechanochemistry of Solid Surfaces", World Scientific Publishing Co., 1994.
  10. G.W. Greenwood, "The Solubility of Gas Bubbles", Journal of Material Science, 4, 320-322, 1969.
  12. Data taken from the Merck Index, 7th edition, Merck & Co., 1960.
  13. Gil A, Chamayou A, Leverd E, Bougaret J, Baron M, Couarraze G (2004). "Evolution of the interaction of a new chemical entity, eflucimibe, with gamma-cyclodextrin during kneading process". Eur. J. Pharm. Sciences, 23, 123-129.
  14. NIST solubility database
  15. ONS Solubility challenge
  16. Solubility of Vanillin in various non-aqueous solvents
  17. O.M.Saether & P. de Caritat (ed.) "Geochemical processes, weathering and groundwater recharge in catchments", Taylor & Francis, Rotterdam, 1997, page 6.

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