Solvation, also sometimes 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.
Distinction between solvation, dissolution and solubility 
By an IUPAC definition, 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 is, in concept, distinct from dissolution and solubility. Dissolution is a kinetic process, and 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.
Liquefaction accompanied by an irreversible chemical change is also distinct from solvation. For example, zinc cannot be solvated by hydrochloric acid, but it can be converted into the soluble salt zinc chloride by a chemical reaction.
Solvents and intermolecular interactions 
Polar solvents are those with a molecular structure that contains dipoles. Such compounds are often found to have a high dielectric constant. The polar molecules of these solvents 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 and creates a solvation shell (or hydration shell in the case of water). Water is the most common and well-studied polar solvent, but others exist, such as acetonitrile, dimethyl sulfoxide, methanol, propylene carbonate, ammonia, ethanol, and acetone. 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 be thermodynamically favored only 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 and molecules 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, whereas 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 ionic radius, 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, it is apparent, 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.
In general, thermodynamic analysis of solutions is done by modeling them as reactions. For example; if you add sodium chloride(s) to water, the salt will dissociate into the ions sodium(+aq) and chloride(-aq). The equilibrium constant for this dissociation can be predicted by the change in Gibb's free energy of this reaction.
Max Born has developed the first quantitative model for solvation of ionic compounds.
See also 
|Look up solvation in Wiktionary, the free dictionary.|
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- Complex (chemistry)
- Solubility equilibrium
- Ideal solution
- Water model
Further reading 
- Dogonadze, Revaz R.; et al. (eds.) (1985-88). The Chemical Physics of Solvation (3 vols. ed.). Amsterdam: Elsevier. ISBN 0-444-42551-9 (part A), ISBN 0-444-42674-4 (part B), ISBN 0-444-42984-0 (Chemistry) Check
- IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "solvation".
- Alireza Mashaghi et al., Hydration strongly affects the molecular and electronic structure of membrane phospholipids. J. Chem. Phys. 136, 114709 (2012) http://jcp.aip.org/resource/1/jcpsa6/v136/i11/p114709_s1
- Mischa Bonn et al., Interfacial Water Facilitates Energy Transfer by Inducing Extended Vibrations in Membrane Lipids, J Phys Chem, 2012 http://pubs.acs.org/doi/abs/10.1021/jp302478a