Nonelectrolyte Solutes

Nonelectrolyte solutes that are not part of the primary solubilized system (solvent-surfactant-additive) can have a significant effect on the solubilizing 

Unlike polar cosolutes with relatively large hydrophobic tails, short-chain alcohols such as ethanol can significantly reduce the solubiUzing power of a surfactant. In the earlier discussion of the effects of such materials on the micellization process, it was shown that the addition of significant quantities of short-chain alcohols, acetone, dioxane, and similar compounds could result in profound changes in the cmc and aggregation number of surfactants, even to the point of completely inhibiting micelle formation. It is understandable, then, that such solutes would also adversely affect the solubilization capacity of a surfactant solution. 

From the above, it seems clear that the effects of an added nonelectrolyte on the solubilizing capacity of a given surfactant system may be quite complex and may not lend itself to easy analysis. It can be assumed, however, that the fundamental relationships that exist between the solutes and the micellization characteristics of the surfactant, in the absence of the solubihzed additive, can be used to good advantage in predicting what may reasonably be expected in the four-component system. 

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McGraw-HiU, New York, 1987. Sandler, S.I., Chemical and Engineeiing Thermodynamics, 2d ed., Wiley, New York, 1989. Smith, J.M., H.C. Van Ness, and M.M. Abbott, Introduction to Chemical Engineeiing Theimodynamics, 5th ed., McGraw-Hill, New York, 1996. Van Ness, H.C., and M.M. Abbott, Classical Theimodynamics of Nonelectrolyte Solutions With Applications to Phase Equi-lihiia, McGraw-Hill, New York, 1982. 

VI. Van Ness, H. C., Classical Thermodynamics of Nonelectrolyte Solutions. Pergamon, Oxford, 1964,... 

Solutions containing strong electrolyte solutes differ from those containing nonelectrolyte solutes in that deviations from Henry s law become important at much lower concentrations for the electrolyte solute than for the nonelectrolyte... 

For a nonelectrolyte solute, it simplifies even further to the expression... 

The solute in an aqueous strong electrolyte solution is present as ions that can conduct electricity through the solvent. The solutes in nonelectrolyte solutions are present as molecules. Only a small fraction of the solute molecules in weak electrolyte solutions are present as ions. 

FIGURE 1.3 In a nonelectrolyte solution, the solute remains as a molecule and does not break up into ions. Methanol, CH)OH, is a nonelectrolyte and is present as intact molecules when it is dissolved in water. 

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. 

It is found empirically and can be justified thermodynamically that the freezing-point depression for an ideal solution is proportional to the molality of the solute. For a nonelectrolyte solution. 

Here, i, the van t Hoff i factor, is determined experimentally. In a very dilute solution (less than about 10 3 mol-I. ), when all ions are independent, i = 2 for MX salts such as NaCl, i = 3 for MX2 salts such as CaCl2, and so on. For dilute nonelectrolyte solutions, i =l. The i factor is so unreliable, however that it is best to confine quantitative calculations of freezing-point depression to nonelectrolyte solutions. Even these solutions must be dilute enough to be approximately ideal. 

In contrast to nonelectrolyte solutions, in the case of electrolyte solutions the col-ligative properties depart appreciably from the values following from the equations above, even in highly dilute electrolyte solutions that otherwise by all means can be regarded as ideal (anomalous colligative properties). 

H.C. Van Ness and M.M. Abbott, Classical Thermodynamics of Nonelectrolytic Solutions, McGraw-Hill Co., New York, NY (1982). 

Clarke, E.C.W. and D.N. Glew (1971), Aqueous nonelectrolyte solutions, Part VIII. Deuterium and hydrogen sulfides solubilities in deuterium oxide and water, Can. J. Chem., 49, 691-698. Corsi, R.L., S. Birkett, H. Melcer, and J. Bell (1995), Control of VOC emissions from sewers A multi-parameter assessment, Water Sci. Tech, 31(7), 147-157. 

Osmotic pressure is a colligative property and is dependent on the number of particles of solute(s) in a solution. The total number of particles of a solute in a solution is the sum of the undissociated molecules and the number of ions into which the molecule dissociates. The number of ions, in turn, depends on the degree of ionization. Thus, a chemical that is highly ionized contributes a greater number of particles to the solution than the same amount of a poorly ionized chemical. When a chemical is a nonelectrolyte such as sucrose or urea, the concentration of the solution depends only on the number of molecules present. The values of the osmotic pressure and other colligative properties are approximately the same for equal concentrations of different nonelectrolyte solutions. 

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... 

Long, F.A. McDevit, W.F. "Activity Coefficients of Nonelectrolyte Solutes in Aqueous Salt Solutions," Chem. Rev.,... 

The similarity of the curves on Figure 1 to those for nonelectrolyte solutions is striking. The dashed line representing ay = x can be called "ideal-solution behavior" for these systems, as it is for nonelectrolytes, but it is realized that a statistical model yielding that result would be more complex for the ionic case. Also the Debye-HUckel effect is a departure from this ideal behavior. Nevertheless, it seems worthwhile to explore the use for these systems of the simple equations for nonelectrolytes. One of the simplest and most successful had its origin in the work of van Laar (15) and has been widely used since. 

Valle-Rie tra Project Evolution in the Chemical Process Industries Van Ness and Abbott Classical Thermodynamics of Nonelectrolyte Solutions with Applications to Phase Equilibria Van Winkle Distillation Volk Applied Statistics for Engineers Walas Reaction Kinetics for Chemical Engineers y ... 

In Chapters 16 and 17, we developed procedures for defining standard states for nonelectrolyte solutes and for determining the numeric values of the corresponding activities and activity coefficients from experimental measurements. The activity of the solute is defined by Equation (16.1) and by either Equation (16.3) or Equation (16.4) for the hypothetical unit mole fraction standard state (X2° = 1) or the hypothetical 1-molal standard state (m = 1), respectively. The activity of the solute is obtained from the activity of the solvent by use of the Gibbs-Duhem equation, as in Section 17.5. When the solute activity is plotted against the appropriate composition variable, the portion of the resulting curve in the dilute region in which the solute follows Henry s law is extrapolated to X2 = 1 or (m2/m°) = 1 to find the standard state. 

As we saw in Section 17.5, the activity coefficient of a nonelectrolyte solute can be calculated from the activity coefficient of the solvent, which, in turn, can be obtained from the measurement of colligative properties such as vapor pressure lowering, freezing point depression, or osmotic pressure. We used the Gibbs-Duhem equation in the form [Equation (17.33)]... 

Solvent extraction rarely involves gases, so that other cases should now be considered. Most liquid organic solutes are completely miscible with, or at least highly soluble in, most organic solvents. The case of a liquid solute that forms a solute-rich liquid phase that contains an appreciable concentration of the solvent is related to the mutual solubility of two solvents, and has been discussed in section 2.2. This leaves solid solutes that are in equilibrium with their saturated solution. It is expedient to discuss organic, nonelectrolytic solutes separately from salts or other ionic solutes. 

Acree W. E. Jr. (1984). Thermodynamic Properties of Nonelectrolyte Solutions. New York Academic Press. 

The square root of the cohesive pressure c as defined in eqn. 3.11 has been termed the solubility parameter 5 by Hildebrand and Scott (1962) because of its value in correlating and predicting the solvency of solvents for nonelectrolyte solutes. Solvency is defined as the ability of solvents to dissolve a compound. A selection of 5-values is given in table 3.10. 

The above equations for e and ky a are for nonelectrolyte solutions. For electrolyte solutions, it is suggested that e and k a be increased by approximately 25%. In electrolyte solutions the bubbles are smaller, the gas holdups are larger, and the interfacial areas larger than in nonelectrolyte solutions. 

Table VII collects the results for all monovalent ion systems for which spectroscopic data are available. Studies of preferential solvation are still at a stage comparable to the establishment of Raoult s and Henry s laws for binary nonelectrolyte solutions. Correlation with thermodynamic data is encouraging for isodielectric solvent systems, but further consideration of the electrostatic terms necessary in the discussion of other systems is required. It is hoped that this present work, which coordinates, correlates, and advances progress made by other workers (7, 18,19, 20, 45, 46, 61, 62, 66, 67, 68), will stimulate systematic experimental investigations of suitable systems by both spectroscopic and thermodynamic methods.

The osmotic coefficient of the binary nonelectrolyte solution was obtained from isopiestic measurements... 

Maron,S.H., Nakajima,N. A theory of the thermodynamic behavior of nonelectrolyte solutions. III. The osmotic pressure of polymer solutions. J. Polymer Sci. 42, 327-340 (1966). 

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