Solvation

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Not to be confused with salvation.
A sodium ion solvated by water molecules.

Solvation, also sometimes called dissolution, is the attraction and association of molecules of a solvent with molecules or ions of a solute. When a solute is soluble in a certain solvent, the solute’s molecules or ions will spread out and become surrounded by solvent molecules. A molecule or ion of solute surrounded by solvent is known as a solvation complex. Solvation is the process of reorganizing solvent and solute molecules into solvation complexes until the solute is distributed evenly within the solvent. Solvation depends on factors such as hydrogen bonding and van der Waals forces. Insoluble solutes prefer to maintain interactions among solute molecules rather than break apart and become solvated by the solvent. Solvation of a solute by water is called hydration.

Distinction between solvation and solubility[edit]

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. In the solvated state, an ion in a solution is surrounded or complexed by solvent molecules. Solvated species can be described by coordination number and the complex stability constants. 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 solubility. Solvation or 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. The typical unit for dissolution rate is mol/s. The units for solubility express a concentration: mass per volume (mg/mL), molarity (mol/L), etc..[2]

Solvents and intermolecular interactions[edit]

Solvation involves different types of intermolecular interactions: hydrogen bonding, ion-dipole interactions, and van der Waals forces (which consist of dipole-dipole, dipole-induced dipole, and induced dipole-induced dipole interactions). Which of these forces are at play depends on the molecular structure and properties of the solvent and solute. The similarity or complementary character of these properties between solvent and solute determines how well a solute can be solvated by a particular solvent.

Nile red at daylight (top row) and UV-light (second row) in different solvents. From left to right: 1. Water, 2. Methanol, 3. Ethanol, 4. Acetonitrile, 5. Dimethylformamide, 6. Acetone, 7. Ethylacetate, 8. Dichlormethane 9. n-Hexane, 10. Methyl-tert-Butylether, 11. Cyclohexane, 12. Toluene. Photographer: Armin Kübelbeck, CC-BY-SA, Wikimedia Commons

Solvent polarity is the most important factor in determining how well it solvates a particular solute. Polar solvents have molecular dipoles, meaning that part of the solvent molecule has more electron density than another part of the molecule. The part with more electron density will experience a partial negative charge while the part with less electron density will experience a partial positive charge. Polar solvent molecules can solvate polar solutes and ions because they can orient the appropriate partially charged portion of the molecule towards the solute through electrostatic attraction. This stabilizes the system and creates a solvation shell (or hydration shell in the case of water) around each particle of solute. The solvent molecules in the immediate vicinity of a solute particle often have a much different ordering than the rest of the solvent, and this area of differently ordered solvent molecules is called the cybotactic region.[3] Water is the most common and well-studied polar solvent, but others exist, such as ethanol, methanol, acetone, acetonitrile, and dimethyl sulfoxide. Polar solvents are often found to have a high dielectric constant, although other solvent scales are also used to classify solvent polarity. Polar solvents can be used to dissolve inorganic or ionic compounds such as salts. The conductivity of a solution depends on the solvation of its ions. Nonpolar solvents cannot solvate ions, and ions will be found as ion pairs.

Hydrogen bonding among solvent and solute molecules depends on the ability of each to accept H-bonds, donate H-bonds, or both. Solvents that can donate H-bonds are referred to as protic, while solvents that do not contain a polarized bond to a hydrogen atom and cannot donate a hydrogen bond are called aprotic. H-bond donor ability is classified on a scale (α).[4] Protic solvents can solvate solutes that can accept hydrogen bonds. Similarly, solvents that can accept a hydrogen bond can solvate H-bond-donating solutes. The hydrogen bond acceptor ability of a solvent is classified on a scale (β).[5] Solvents such as water can both donate and accept hydrogen bonds, making them excellent at solvating solutes that can donate or accept (or both) H-bonds.

Some chemical compounds experience solvatochromism, which is a change in color due to solvent polarity. This phenomenon illustrates how different solvents interact differently with the same solute. Other solvent effects include conformational or isomeric preferences and changes in the acidity of a solute.

Solvation energy & thermodynamic considerations[edit]

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. It is important to remember, however, that 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 ethalpic 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.[3]

Solvation of a solute by solvent

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. The sum of the enthalpy and entropy changes throughout these steps is called the solvation energy.

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.[6]

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 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.[3]

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.

Macromolecules and assemblies[edit]

Solvation (specifically, hydration) is important for many biological structures and processes. For instance, protein folding occurs spontaneously because of a favorable change in the interactions between the protein and the surrounding water molecules. Proteins are stabilized by 5-10 kcal/mol when folded due to a combination of solvent and hydrogen bonding effects.[7] Minimizing the number of hydrophobic side-chains exposed to water by burying them in the center of a folded protein is a driving force related to solvation. Solvation can also drive host-guest complexation. Many host molecules have a hydrophobic pore that readily encapsulates a hydrophobic guest. These interactions can be used in applications such as drug delivery, such that a hydrophobic drug molecule can be delivered in a biological system without needing to covalently modify the drug in order to solubilize it. Binding constants for host-guest complexes depend on the polarity of the solvent.[8]

Hydration affects electronic and vibrational properties of biomolecules.[9][10]

See also[edit]

References[edit]

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "solvation".
  2. ^ "Solubility - Common Measuring Units" "http://science.jrank.org/pages/6279/Solubility-Common-measuring-units.html">Solubility - Common Measuring Units
  3. ^ a b c Eric V. Anslyn; Dennis A. Dougherty (2006). Modern Physical Organic Chemistry. University Science Books. ISBN 978-1-891389-31-3.
  4. ^ Taft, R. W. & Kamlet, M. J. "The solvatochromic comparison method. 2. The .alpha.-scale of solvent hydrogen-bond donor (HBD) acidities" J. Am. Chem. Soc., 1976, 98 (10), pp 2886–2894
  5. ^ Taft, R. W. & Kamlet, M. J. "The solvatochromic comparison method. 1. The .beta.-scale of solvent hydrogen-bond acceptor (HBA) basicities" J. Am. Chem. Soc., 1976, 98 (2), pp 377-383
  6. ^ V. Gutmann (1976). "Solvent effects on the reactivities of organometallic compounds". Coord. Chem. Rev. 18 (2): 225. doi:10.1016/S0010-8545(00)82045-7
  7. ^ Pace CN, Shirley BA, McNutt M, Gajiwala K (January 1996). "Forces contributing to the conformational stability of proteins". FASEB Journal. 10 (1): 75–83. PMID 8566551
  8. ^ Steed, J. W. and Atwood, J. L. (2013) Supramolecular Chemistry. 2nd ed. Wiley. ISBN 1118681509, 9781118681503.
  9. ^ Alireza Mashaghi et al., Hydration strongly affects the molecular and electronic structure of membrane phospholipids. J. Chem. Phys. 136, 114709 (2012) doi:10.1063/1.3694280
  10. ^ Mischa Bonn et al., Interfacial Water Facilitates Energy Transfer by Inducing Extended Vibrations in Membrane Lipids, J Phys Chem, 2012 doi:10.1021/jp302478a

Further reading[edit]

  • 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)
  • Jiang D., Urakawa A., Yulikov M., Mallat T., Jeschke G. & Baiker A., 2009, "Size selectivity of a copper metal-organic framework and origin of catalytic activity in epoxide alcoholysis," Chemistry 15(45):12255-62, DOI 10.1002/chem.200901510. [One example of a solvated MOF, where partial dissolution is described.]

External links[edit]