Nonelectrolytes, interaction with ions

A hypothetical solution that obeys Raoult s law exactly at all concentrations is called an ideal solution. In an ideal solution, the interactions between solute and solvent molecules are the same as the interactions between solvent molecules in the pure state and between solute molecules in the pure state. Consequently, the solute molecules mingle freely with the solvent molecules. That is, in an ideal solution, the enthalpy of solution is zero. Solutes that form nearly ideal solutions are often similar in composition and structure to the solvent molecules. For instance, methylbenzene (toluene), C6H5CH, forms nearly ideal solutions with benzene, C6H6. Real solutions do not obey Raoult s law at all concentrations but the lower the solute concentration, the more closely they resemble ideal solutions. Raoult s law is another example of a limiting law (Section 4.4), which in this case becomes increasingly valid as the concentration of the solute approaches zero. A solution that does not obey Raoult s law at a particular solute concentration is called a nonideal solution. Real solutions are approximately ideal at solute concentrations below about 0.1 M for nonelectrolyte solutions and 0.01 M for electrolyte solutions. The greater departure from ideality in electrolyte solutions arises from the interactions between ions, which occur over a long distance and hence have a pronounced effect. Unless stated otherwise, we shall assume that all the solutions that we meet are ideal. 

An understanding of equilibrium phenomena in naturally occurring aqueous systems must, in the final analysis, involve understanding the interaction between solutes and water, both in bulk and in interfacial systems. To achieve this goal, it is reasonable to attempt to describe the structure of water, and when and if this can be achieved, to proceed to the problems of water structure in aqueous solutions and solvent-solute interactions for both electrolytes and nonelectrolytes. This paper is particularly concerned with two aspects of these problems—current views of the structure of water and solute-solvent interactions (primarily ion hydration). It is not possible here to give an exhaustive account of all the current structural models of water instead, we shall describe only those which may concern the nature of some reported thermal anomalies in the properties of water and aqueous solutions. Hence, the discussion begins with a brief presentation of these anomalies, followed by a review of current water structure models, and a discussion of some properties of aqueous electrolyte solutions. Finally, solute-solvent interactions in such solutions are discussed in terms of our present understanding of the structural properties of water. 

In electrolytic solutions we were concerned with electrostatic interactions between ions in the solution and with the solvent (water). In solutions of nonelectrolytes we will be concerned with molecule-solvent interactions due to electrostatic forces, dispersion forces, and chemical forces. 

Any substance whose aqueous solution contains ions is called an electrolyte. Any substance that forms a solution containing no ions is a nonelectrolyte. Electrolytes that are present in solution entirely as ions are strong electrolytes, whereas those that are present partly as ions and partly as molecules are weak electrolytes. Ionic compounds dissociate into ions when they dissolve, and they are strong electrolytes. The solubility of ionic substances is made possible by solvation, the interaction of ions with polar solvent molecules. Most molecular compounds are nonelectrolytes, although some are weak electrolytes, and a few are strong electrolytes. When representing the ionization of a weak electrolyte in solution, half-arrows in both directions are used, indicating that the forward and reverse reactions can achieve a chemical balance called a chemical equilibrium. 

Nonelectrolytes in electrolyte solutions interact with both the free solvent molecules and with the solvated ions and ion associates. A general approach to the study of such interactions is in terms of the solubility of the nonelectrolyte, subsalpt N, in the solvent in the presence of an electrolyte, compared with the solubility in its absence. 

It is interesting to note that the molecule-ion interaction contribution in equation (5) is consistent with the well-known Setschenow equation. The Setschenow equation is used to represent the salting-out effect of salts on molecular nonelectrolyte solutes, when the solubilities of the latter are small (Gordon, (15)). The Setschenow equation is... 

Some polar nonelectrolytes show nonlinear Setschenow plots, and this behavior has been attributed to direct chemical interactions (Gordon, 1975). If a nonelectrolyte forms a complex ion with one of the ions of the salt, the solubility of a poorly soluble solute may in-... 

The identical slopes observed in Figures 8-10 suggest that at low concentrations Ben = Bne this means that the interactions of a single ion with the surrounding nonelectrolytes is the same as that of a single nonelectrolyte with the surrounding ions provided only pair interactions are involved. When triplet and higher interactions become important, this identity is lost. 

Dispersion forces give rise to an interaction energy in which the potential energy of interaction varies as r , where r is the distance between the centers of the two substances interacting. Thus, the equation for the dispersive energy of interaction may be written as A/r , where A is a constant independent of r. The rapid decrease of such forces with increase of distance from the origin makes it unnecessary to consider dispersion interactions outside the primary solvation shell by then, they have already decreased to an extent that they no longer warrant consideration. Inside the primary hydration sheath, the dispersion interaction can be treated in the same way as the ion-dipole interaction. That is, in the replacement of a water molecule by a nonelectrolyte molecule, one must take into account not only the difference in ion-dipole... 

Colligative properties are related to the number of dissolved solute particles, not their chemical nature. Compared with the pure solvent, a solution of a nonvolatile nonelectrolyte has a lower vapor pressure (Raoult s law), an elevated boiling point, a depressed freezing point, and an osmotic pressure. Colligative properties can be used to determine the solute molar mass. When solute and solvent are volatile, the vapor pressure of each is lowered by the presence of the other. The vapor pressure of the more volatile component is always higher. Electrolyte solutions exhibit nonideal behavior because ionic interactions reduce the effective concentration of the ions. 

Nonelectrolyte G mcxlels only account for the short-range interaction among non-charged molecules (—One widely used G model is the Non-Random-Two-Liquid (NRTL) theory developed in 1968. To extend this to electrolyte solutions, it was combined with either the DH or the MSA theory to explicitly account for the Coulomb forces among the ions. Examples for electrolyte models are the electrolyte NRTL (eNRTL) [4] or the Pitzer model [5] which both include the Debye-Hiickel theory. Nasirzadeh et al. [6] used a MSA-NRTL model [7] (combination of NRTL with MSA) as well as an extended Pitzer model of Archer [8] which are excellent models for the description of activity coefficients in electrolyte solutions. Examples for electrolyte G models which were applied to solutions with more than one solvent or more than one solute are a modified Pitzer approach by Ye et al. [9] or the MSA-NRTL by Papaiconomou et al. [7]. However, both groups applied ternary mixture parameters to correlate activity coefficients. Salimi et al. [10] defined concentration-dependent and salt-dependent ion parameters which allows for correlations only but not for predictions or extrapolations. 

Here eiec is the molar electrostriction by the electrolyte, and JCjw is the isothermal compressibility of the water. The direct interactions of the ions of the electrolyte with molecules of the nonelectrolyte are ignored in this approach. Actually, values of predicted by Eq. 2 are several-fold larger than the experimental values [2, 3]. 

We now leave pure materials and the limited but important changes they can undergo and examine mixtures. We shall consider only homogeneous mixtures, or solutions, in which the composition is uniform however small the sample. The component in smaller abundance is called the solute and that in larger abimdance is the solvent. These terms, however, are normally but not invariably reserved for solids dissolved in Kquids one liquid mixed with another is normally called simply a mixture of the two liquids. In this chapter we consider mainly nonelectrolyte solutions, where the solute is not present as ions. Examples are sucrose dissolved in water, sulfur dissolved in carbon disulfide, and a mixture of ethanol and water. Although we also consider some of the special problems of electrolyte solutions, in which the solute consists of ions that interact strongly with one another, we defer a full study until Chapter 5. The measures of concentration commonly encoimtered in physical chemistry are reviewed in Further information 3.2. 

Both nonelectrolyte and electrolyte additives cause dehydration of nonionic micellar aggregates in aqueous solvents. The dehydration effect of these additives facilitates the expansion of the hydrophobic microenvironment of nonionic micelles and consequently increases the stability of micelles. This interaction is similar to what is known as salting-out effect of hydrophilic ions. Ions with a large charge density increase the water-structure and hence, the increase in the concentration of water-structure-forming hydrophilic ions is expected to cause nonlinear decrease in CMC of nonionic surfactants. 

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