Editing Solvation (section)

== Solvation energy and thermodynamic considerations ==
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 energy of the system decreases. A negative Gibbs energy indicates a spontaneous process but does not provide information about the rate of dissolution.

Solvation involves multiple steps with different energy consequences.  First, a cavity must form in the solvent to make space for a solute. This is both entropically and enthalpically unfavorable, as solvent ordering increases and solvent-solvent interactions decrease. Stronger interactions among solvent molecules leads to a greater enthalpic penalty for cavity formation. Next, a particle of solute must separate from the bulk. This is enthalpically unfavorable since solute-solute interactions decrease, but when the solute particle enters the cavity, the resulting solvent-solute interactions are enthalpically favorable. Finally, as solute mixes into solvent, there is an entropy gain.<ref name="Anslyn & Dougherty" />

[[File:Effect of solvent on solubility.png|600px|Solvation of a solute by solvent]]

The [[enthalpy of solution]] is the solution enthalpy minus the enthalpy of the separate systems, whereas the entropy of solution is the corresponding difference in [[entropy]]. The solvation energy (change in [[Gibbs free energy]]) is the change in enthalpy minus the product of temperature (in [[Kelvin]]) times the change in entropy. Gases have a negative entropy of solution, due to the decrease in gaseous volume as gas dissolves. Since their enthalpy of solution does not decrease too much with temperature, and their entropy of solution is negative and does not vary appreciably with temperature, most gases are less soluble at higher temperatures.

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 number|negative]] value for the enthalpy change of solution corresponds to an ion that is likely to dissolve, whereas a high [[positive number|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 number]]s.<ref>{{cite journal | author = Gutmann V | year = 1976 | title = Solvent effects on the reactivities of organometallic compounds | journal = Coord. Chem. Rev. | volume = 18 | issue = 2| page = 225 | doi = 10.1016/S0010-8545(00)82045-7 }}</ref>

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 [[lanthanide]]s and [[actinide]]s, which are readily hydrolyzed to form insoluble (hydrous) oxides. As these are solids, it is apparent that they are not solvated.

Strong solvent–solute interactions make the process of solvation more favorable. One way to compare how favorable the dissolution of a solute is in different solvents is to consider the free energy of transfer. The free energy of transfer quantifies the free energy difference between dilute solutions of a solute in two different solvents. This value essentially allows for comparison of solvation energies without including solute-solute interactions.<ref name="Anslyn & Dougherty" />

In general, thermodynamic analysis of solutions is done by modeling them as reactions. For example, if you add sodium chloride 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 Gibbs energy of this reaction.

The [[Born equation]] is used to estimate Gibbs free energy of solvation of a gaseous ion.

Recent simulation studies have shown that the variation in solvation energy between the ions and the surrounding water molecules underlies the mechanism of the [[Hofmeister series]].<ref>{{cite journal |author=M. Andreev|author2=A. Chremos |author3=J. de Pablo|author4=J. F. Douglas |title=Coarse-Grained Model of the Dynamics of Electrolyte Solutions|journal=J. Phys. Chem. B |volume=121 |issue=34 |pages=8195–8202|year=2017|doi=10.1021/acs.jpcb.7b04297|pmid=28816050 }}</ref><ref name="article-1">{{cite journal |author=M. Andreev|author2=J. de Pablo |author3=A. Chremos|author4=J. F. Douglas |title=Influence of Ion Solvation on the Properties of Electrolyte Solutions|journal=J. Phys. Chem. B |volume=122 |issue=14 |pages=4029–4034|year=2018|doi=10.1021/acs.jpcb.8b00518|pmid=29611710 }}</ref>