Solubility is the property of a solid, liquid, or gaseous chemical substance called solute to dissolve in a liquid solvent to form a homogeneous solution. The solubility of a substance strongly depends on the used solvent as well as on temperature and pressure. The pressure also affects the solution whether it is gas or liquid, like temperature. So, in definition of solubility we always mention the pressure and temperature "fixed". The extent of the solubility of a substance in a specific solvent is measured as the saturation concentration where adding more solute does not increase the concentration of the solution.
The solvent is generally a liquid, which can be a pure substance or a mixture. 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 ) 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.
Under certain conditions the equilibrium solubility can be exceeded to give a so-called supersaturated solution, which is metastable.
The solvent is generally a liquid, which can be a pure substance or a mixture. 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 ) 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.
Under certain conditions the equilibrium solubility can be exceeded to give a so-called supersaturated solution, which is metastable.
Factors affecting solubility:
Solubility is defined for specific 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 in another is determined by the balance of intermolecular forces 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-forming anions (ligands) 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 solutions. The last two effects can be quantified using the equation for solubility equilibrium.
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.
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.[citation needed] 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. 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.
For condensed phases (solids and liquids), the pressure dependence of solubility is typically weak and usually neglected in practice.
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. 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.
For condensed phases (solids and liquids), the pressure dependence of solubility is typically weak and usually neglected in practice.
Polarity:
A popular aphorism used for predicting solubility is "like dissolves like". 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.
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.
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.
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.
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.
Solubility of organic compounds:
The principle outlined above under polarity, that like dissolves like, is the usual guide to solubility with organic systems. For example, petroleum jelly will dissolve in gasoline because both petroleum jelly and gasoline are hydrocarbons. 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 extraction with water.
place in a space between the lattice points (an interstitial position, for example: carbon in iron).
In microelectronic fabrication, solid solubility refers to the maximum concentration of impurities
one can place into the substrate.
References:
1. ^ Yuen, C. (2003), Element, Compound and Mixture
2. ^ Clugston M. and Fleming R. (2000), p.108
3. ^ Cancerweb.ncl.ac.uk: 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 Books.Google.com.
5. ^ a b 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.
7. ^ Data taken from the Handbook of Chemistry and Physics, 27th edition, Chemical Rubber Publishing Co., Cleveland, Ohio, 1943.
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. ^ Kenneth J. Williamson, Macroscale and Microscale Organic Experiments, p40, 2nd edition, D. C, Heath, Lexington, Mass., 1994.
1. ^ Yuen, C. (2003), Element, Compound and Mixture
2. ^ Clugston M. and Fleming R. (2000), p.108
3. ^ Cancerweb.ncl.ac.uk: 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 Books.Google.com.
5. ^ a b 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.
7. ^ Data taken from the Handbook of Chemistry and Physics, 27th edition, Chemical Rubber Publishing Co., Cleveland, Ohio, 1943.
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. ^ Kenneth J. Williamson, Macroscale and Microscale Organic Experiments, p40, 2nd edition, D. C, Heath, Lexington, Mass., 1994.
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