Solvent—solute interactions

Because of favorable dipole-dipole attractions between solvent molecules and solute molecules, polar liquids tend to dissolve in polar solvents. Water is both polar and able to form hydrogen bonds. (Section 11.2) Thus, polar molecules, especially those that can form hydrogen bonds with water molecules, tend to be soluble in water. For example, acetone, a polar molecule with the structural formula shown in the margin, mixes in all proportions with water. Acetone has a strongly polar C = O bond and pairs of nonbonding electrons on the O atom that can form hydrogen bonds with water. 

Pairs of liquids that mix in all proportions, such as acetone and water, are miscible, whereas those that do not dissolve in one another are immiscible. Gasoline, which is a mixture of hydrocarbons, is immiscible with water. Hydrocarbons are nonpolar substances because of several factors The C—C bonds are nonpolar, the C—H bonds are nearly nonpolar, and the molecules are symmetrical enough to cancel much of the weak C— H bond dipoles. The attraction between the polar water molecules and the nonpolar hydrocarbon molecules is not sufficiently strong to allow the formation of a solution. Nonpolar liquids tend to be insoluble in polar liquids, as FIGURE 13.10 shows for hexane (CgHj4) and water. 

A FIGURE 13.10 Hexane is immiscible with water. Hexane is the tep layer because it is less dense than water. 

Hydrogen bond between ethanol molecule and water molecule 

Notice in Table 13.2 that the number of carbon atoms in an alcohol affects its solubility in water. As this number increases, the polar OH group becomes an ever smaller part of the molecule, and the molecule behaves more like a hydrocarbon. The solubility of the alcohol in water decreases correspondingly. On the other hand, the solubility of alcohols in a nonpolar solvent like hexane (CgHi4) increases as the nonpolar hydrocarbon chain lengthens. 

Cyclohexane, C Hq, which has no polar OH groups, is essentially insoluble in water 

OH groups enhance the aqueous solubility because of their ability to hydrogen bond with H2O. 

Stokes-Einstein equation with some other theoretically rationalized value. The values calculated from Eq. 6.2-45, shown in the fourth column, do have the courtesy to remain positive, but they are far from being integers. I am always unsure how a cesium ion can react with half a molecule of water. In addition, the hydration numbers found from diffusion show little relation with those calculated from values from the other types of experiments shown in Table 6.2-1. These ideas have only qualitative value. 

In every case, diffusion is about mixing. In almost every case in this book, we are interested in what happens when two miscible solutions are placed next to each other and then allowed to mix without flow as the result of molecular motion. The speed of this spontaneous mixing is described by diffusion. This diffusion is a consequence of free energy decreases, of the second law of thermodynamics. 

In some cases, diffusion occurs much more slowly than expected. This most commonly occurs near a phase boundary where the solution is supersaturated. An example is diffusion in supersaturated solutions of sugar in water, shown in Fig. 6.3-1. Slow diffusion also occurs in solutions near to a consolute point where two liquids first become miscible. Examples of diffusion near consolute points are shown in Fig. 6.3-2. Other related cases occur when an initially homogeneous solution is suddenly quenched to cause a phase separation. This quenching is commonly effected by abruptly lowering the temperature. The phase separation then occurs very rapidly, at a rate proportional to the diffusion. 

Each of these cases involves mass transfer driven by changes in free energy, or more exactly, by gradients in chemical potential. Their description requires major changes in 

Pick s law, including the effects of higher order terms in the gradients. In particular, the form postulated for the flux /i is 

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The solute-solvent interaction in equation A2.4.19 is a measure of the solvation energy of the solute species at infinite dilution. The basic model for ionic hydration is shown in figure A2.4.3 [5] there is an iimer hydration sheath of water molecules whose orientation is essentially detemiined entirely by the field due to the central ion. The number of water molecules in this iimer sheath depends on the size and chemistry of the central ion ... 

Specific solute-solvent interactions involving the first solvation shell only can be treated in detail by discrete solvent models. The various approaches like point charge models, siipennoleciilar calculations, quantum theories of reactions in solution, and their implementations in Monte Carlo methods and molecular dynamics simulations like the Car-Parrinello method are discussed elsewhere in this encyclopedia. Here only some points will be briefly mentioned that seem of relevance for later sections. 

Considering, for simplicity, only electrostatic interactions, one may write the solute-solvent interaction temi of the Hamiltonian for a solute molecule surrounded by S solvent molecules as... 

Kramers solution of the barrier crossing problem [45] is discussed at length in chapter A3.8 dealing with condensed-phase reaction dynamics. As the starting point to derive its simplest version one may use the Langevin equation, a stochastic differential equation for the time evolution of a slow variable, the reaction coordinate r, subject to a rapidly statistically fluctuating force F caused by microscopic solute-solvent interactions under the influence of an external force field generated by the PES F for the reaction... 

RRKM fit to microcanonical rate constants of isolated tran.s-stilbene and the solid curve a fit that uses a reaction barrier height reduced by solute-solvent interaction [46],... 

Aguilar M A and Olivares del Valle F J 1989 Solute-solvent interactions. A simple procedure for constructing the solvent capacity for retaining a molecular solute Ohem. Rhys. 129 439-50... 

Schroeder J 1996 The role of solute-solvent interactions in the dynamics of unimolecular reactions in compressed solvents J. Phys. Condens. Matters 9379... 

The variation of chemical shifts as a function of dilution could be accounted for only qualitatively (235) because of the large diversity of solute-solvent interactions resulting from the nature and the shape of the solvent molecule (Table 1-34). 

The nature of solute-solute and solute-solvent interactions is dependent on the solvent environment. Solvent influences the hydrogen-bonding pattern, solute surface area, and hydrophilic and hydrophobic group exposures. 

Q are the absorbance and wavenumber, respectively, at the peak (center) of the band, p is the wavenumber, and y is the half width of the band at half height. Liquid band positions ate usually shifted slightly downward from vapor positions. Both band positions and widths of solute spectra are affected by solute—solvent interactions. Spectra of soHd-phase samples are similar to those of Hquids, but intermolecular interactions in soHds can be nonisotropic. In spectra of crystalline samples, vibrational bands tend to be sharper and may spHt in two, and new bands may also appear. If polarized infrared radiation is used, both crystalline samples and stressed amorphous samples (such as a stretched polymer film) show directional effects (28,29). 

The solvophobic model of Hquid-phase nonideaHty takes into account solute—solvent interactions on the molecular level. In this view, all dissolved molecules expose microsurface area to the surrounding solvent and are acted on by the so-called solvophobic forces (41). These forces, which involve both enthalpy and entropy effects, are described generally by a branch of solution thermodynamics known as solvophobic theory. This general solution interaction approach takes into account the effect of the solvent on partitioning by considering two hypothetical steps. Eirst, cavities in the solvent must be created to contain the partitioned species. Second, the partitioned species is placed in the cavities, where interactions can occur with the surrounding solvent. The idea of solvophobic forces has been used to estimate such diverse physical properties as absorbabiHty, Henry s constant, and aqueous solubiHty (41—44). A principal drawback is calculational complexity and difficulty of finding values for the model input parameters. 

As long as the normalization condition given by Eq. (5) is satisfied, an arbitrary offset constant may be added to W(X) without affecting averages in Eq. (3). The absolute value of the PMF is thus unimportant. For convenience, it is possible to choose the value of the free energy W(X) relative to a reference system from which the solute-solvent interactions are absent. The free energy W(X) may thus be expressed as... 

The electrostatic free energy contribution in Eq. (14) may be expressed as a thennody-namic integration corresponding to a reversible process between two states of the system no solute-solvent electrostatic interactions (X = 0) and full electrostatic solute-solvent interactions (X = 1). The electrostatic free energy has a particularly simple form if the thermodynamic parameter X corresponds to a scaling of the solute charges, i.e., (X,... 

Solvent effects on chemical equilibria and reactions have been an important issue in physical organic chemistry. Several empirical relationships have been proposed to characterize systematically the various types of properties in protic and aprotic solvents. One of the simplest models is the continuum reaction field characterized by the dielectric constant, e, of the solvent, which is still widely used. Taft and coworkers [30] presented more sophisticated solvent parameters that can take solute-solvent hydrogen bonding and polarity into account. Although this parameter has been successfully applied to rationalize experimentally observed solvent effects, it seems still far from satisfactory to interpret solvent effects on the basis of microscopic infomation of the solute-solvent interaction and solvation free energy. 

J. F. Coetzee and C. D. Ritchie, Solute-Solvent Interactions, Marcel Dekker, New kbtk, 1969. 

Because the key operation in studying solvent effects on rates is to vary the solvent, evidently the nature of the solvation shell will vary as the solvent is changed. A distinction is often made between general and specific solvent effects, general effects being associated (by hypothesis) with some appropriate physical property such as dielectric constant, and specific effects with particular solute-solvent interactions in the solvation shell. In this context the idea of preferential solvation (or selective solvation) is often invoked. If a reaction is studied in a mixed solvent. 

The preceding empirical measures have taken chemical reactions as model processes. Now we consider a different class of model process, namely, a transition from one energy level to another within a molecule. The various forms of spectroscopy allow us to observe these transitions thus, electronic transitions give rise to ultraviolet—visible absorption spectra and fluorescence spectra. Because of solute-solvent interactions, the electronic energy levels of a solute are influenced by the solvent in which it is dissolved therefore, the absorption and fluorescence spectra contain information about the solute-solvent interactions. A change in electronic absorption spectrum caused by a change in the solvent is called solvatochromism. 

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