Solvation

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 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

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: 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 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 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