Solvation

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Solvation, commonly called dissolution, is the process of attraction and association of molecules of a solvent with molecules or ions of a solute. As ions dissolve in a solvent they spread out and become surrounded by solvent molecules. The bigger the ion, the more solvent molecules are able to surround it and the more it becomes solvated.

Distinction between solvation, dissolution and solubility

By an IUPAC definition^[1], solvation is an interaction of a solute with the solvent which leads to stabilization of the solute species in the solution. One may also refer to the solvated state, whereby an ion in a solution is complexed by solvent molecules. The concept of the solvation interaction can also be applied to an insoluble material, for example, solvation of functional groups on a surface of ion-exchange resin.

Solvation should be conceptually separated from dissolution and solubility. Dissolution is a kinetic process and it is quantified by its rate. Solubility quantifies the dynamic equilibrium state achieved when the rate of dissolution equals the rate of precipitation.

The consideration of the units makes the distinction clearer. Complexation can be described by coordination number and the complex stability constants. The typical unit for dissolution rate is mol/s. The unit for solubility can be mol/kg.

Solvents and intermolecular interactions

Solvents that are Polar are ones with a molecular structure that contains polar bonds. Such compounds are often found to have a high dielectric constant. Examples of polar solvents include water and acetonitrile. These polar molecules can solvate ions because they can orient the appropriate partially charged portion of the molecule towards the ion in response to electrostatic attraction. This stabilizes the system. Water represents the most common and well-studied polar solvent, but others exist, such as acetonitrile, dimethyl sulfoxide, methanol, propylene carbonate, ammonia, ethanol, and acetone, among others. These solvents can be used to dissolve inorganic compounds such as salts.

Solvation involves different types of intermolecular interactions: hydrogen bonding, ion-dipole and dipole-dipole attractions or van der Waals forces. The hydrogen bonding, ion-dipole, and dipole-dipole interactions occur only in polar solvents. Ion-ion interactions occur only in ionic solvents. The solvation process will only be thermodynamically favored if the overall Gibbs energy of the solution is decreased compared to the Gibbs energy of the separated solvent and solid (or gas or liquid). This means that the change in enthalpy minus the change in entropy (multiplied by the absolute temperature) is a negative value, or that the Gibbs free energy of the system decreases.

Thermodynamic considerations

For solvation to occur, energy is required to release individual ions from the crystal lattices in which they are present. This is necessary to break the attractions the ions have with each other and is equal to the solid's lattice free energy (the energy released at the formation of the lattice as the ions bonded with each other). The energy for this comes from the energy released when ions of the lattice associate with molecules of the solvent. Energy released in this form is called the free energy of solvation.

The enthalpy of solution is the solution enthalpy minus the enthalpy of the separate systems, while the entropy is the corresponding difference in entropy. Most gases have a negative enthalpy of solution. A negative enthalpy of solution means that the solute is less soluble at high temperatures.

Although early thinking was that a higher ratio of a cation's ion charge to the size, or the charge density, resulted in more solvation, this does not stand up to scrutiny for ions like Iron(III) or lanthanides and actinides, which are readily hydrolyzed to form insoluble (hydrous)oxides. As solids, these are obviously not solvated.

Enthalpy of solvation can help explain why solvation occurs with some ionic lattices but not with others. The difference in energy between that which is necessary to release an ion from its lattice and the energy given off when it combines with a solvent molecule is called the enthalpy change of solution. A negative value for the enthalpy change of solution corresponds to an ion that is likely to dissolve, whereas a high positive value means that solvation will not occur. It is possible that an ion will dissolve even if it has a positive enthalpy value. The extra energy required comes from the increase in entropy that results when the ion dissolves. The introduction of entropy makes it harder to determine by calculation alone whether a substance will dissolve or not. A quantitative measure for solvation power of solvents is given by donor numbers.

Note that solvation does not mean a reaction takes place. Adding NaCl(s) to water, for example, will only create a solution of sodium and chloride ions; you would only have solvation of the salt's ions. Adding the weak base ammonia to water, on the other hand, would be a reaction.

See also

• Complex (chemistry)
• Saturation
• Solubility
• Solubility equilibrium
• Solute
• Solution
• Solvent
• Supersaturation

Further reading

• Dogonadze, Revaz R.; et al. (eds.) (1985-88). The Chemical Physics of Solvation, 3 vols., Amsterdam: Elsevier. ISBN 0-444-42551-9 (part A), ISBN 0-444-42674-4 (part B), ISBN 0-444-42984-0 (part C). 

References

v • d • e Articles related to solutions
Solution	Ideal solution • Aqueous solution • Solid solution • Flory-Huggins • Mixture • Suspension (chemistry) • Colloid • Phase diagram • Eutectic point • Alloy
Concentration	Saturation (chemistry) • Supersaturation • Molar solution • Percentage solution
Solubility	Solubility equilibrium • Total dissolved solids • Dissolve • Solvation • Enthalpy change of solution • Lattice energy • Henry's law • Solubility table (data) • Solubility chart
Solvent	(category) • Acid dissociation constant • Protic solvent • Inorganic nonaqueous solvent • Solvation • Solvation shell • List of boiling and freezing information of solvents Partition coefficient • Polarity • Hydrophobe • Hydrophile • Lipophilic • Amphiphile

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