Properties of Strong Electrolyte Solutions

To calculate the colligative properties of strong electrolyte solutions, we incorporate the van t Hoff factor into the equation  

For boiling point elevation For freezing point depression For osmotic pressure  

If strong electrolyte solutions behaved ideally, the factor i would be the amount (mol) of particles in solution divided by the amount (mol) of dissolved solute that is, i would be 2 for NaCl, 3 for Mg(N03)2, and so forth. Careful experiment shows, however, that most strong electrolyte solutions are not ideal. For example, comparing the boiling point elevation for 0.050 m NaQ solution with that for 0.050 m glucose solution gives a factor i of 1.9, not 2.0  

SAMPLE PROBLEM 13.9 Depicting a Solution to Find Its Colligative Properties 

Problem A 0.952-g sample of magnesium chloride is dissolved in 100. g of water. 

Solution (a) The formula is MgCl2 only scene A has 1 Mg for every 2 Cl . 

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Calculation of the Thermodynamic Properties of Strong Electrolyte Solutes The Debye-Hiickel Theory... 

Perry, R. L., H. Cabezas, and J. P. O Connell. 1988. Fluctuation thermodynamic properties of strong electrolyte-solutions. 63, 189. 

For colligative properties of strong electrolyte solutions, the solute formula tells us the number of particles. For instance, the boiling point elevation iAT ) of 0.050 m NaCl should be 2 X of 0.050 m glucose (C(5Hj20(5), because NaCl dissociates into two particles per formula unit. Thus, we use a multiplying factor called the van t Hoff factor (i), named after the Dutch chemist Jacobus van t Hoff (1852-1911) ... 

The beginning of the twentieth century also marked a continuation of studies of the structure and properties of electrolyte solution and of the electrode-electrolyte interface. In 1907, Gilbert Newton Lewis (1875-1946) introduced the notion of thermodynamic activity, which proved to be extremally valuable for the description of properties of solutions of strong electrolytes. In 1923, Peter Debye (1884-1966 Nobel prize, 1936) and Erich Hiickel (1896-1981) developed their theory of strong electrolyte solutions, which for the first time allowed calculation of a hitherto purely empiric parameter—the mean activity coefficients of ions in solutions. 

Bromley, L. A., "Thermodynamic Properties of Strong Electrolytes in Aqueous Solutions," AIChE J., 1973, 19, 313. 

Additives may also be incorporated into the electrolyte solution to enhance selectivity, which expresses the ability of the separation method to distinguish analytes from each other. Selectivity in CZE is based on differences in the electrophoretic mobility of the analytes, which depends on their effective charge-to-hydrodynamic radius ratio. This implies that selectivity is strongly affected by the pH of the electrolyte solution, which may influence sample ionization, and by any variation of physicochemical property of the electrolyte solution that influences the electrophoretic mobility (such as temperature, for example) [144] or interactions of the analytes with the components of the electrolyte solution which may affect their charge and/or hydrodynamic radius. 

Bromley, L.A. 1973. Thermodynamic properties of strong electrolytes in aqueous solutions. AIChE J. 19 (2) 313-320. 

Unlike weak electrolytes, solutions of strong ones have a far higher specific conductance the rise of the latter with rising concentration is also much more rapid. There is another difference the anomalies ascertained in the colligative properties of strong electrolytes cannot be ascribed to partial dissociation of molecules to ions as in the case of weak electrolytes. 

Thermal Properties of Strong Electrolytes.—According to equation (42) the free energy of an ionic solution may be expressed in the form... 

The colligative properties of an electrolyte solution can be used to determine percent dissociation. Percent dissociation is the percentage of dissolved molecules (or formula units, in the case of an ionic compound) that separate into ions in solution. For a strong electrolyte such as NaCl, there should be complete, or 100 percent, dissociation. However, the data in Table 13.4 indicate that this is not necessarily the case. An experimentally determined van t Hoff factor smaller than the corresponding calculated value indicates less than 100 percent dissociation. As the experimentally determined van t Hoff factors for NaCl indicate, dissociation of a strong electrolyte is more complete at lower concentration. The percent ionization of a weak electrolyte, such as a weak acid, also depends on the concentration of the solution. 

Not only the conductance but also the colligative properties of strong electrolytes show deviations from the values to be expected on the basis of complete ionization. The freezing-point depression of, for example, a NaCl solution is less than we would expect for 2 moles of ions per mole of NaCl the van t Hoff factor approaches 2 only in very dilute solutions (Table 12.2, page 217). These diminutions in the colliga-... 

The Arrhenius theory of electrolytic solutions was quantitatively restored by explaining the properties of strong electrolytes without activity coefficients on basis of ionic solvation and incomplete dissociation [263-266],... 

Until now, we have discussed thermodynamic properties of strong electrolytes in water. In contrast to these systems, weak electrolytes do not fully dissociate in solution. Since in these cases fewer ions are available in solution, the MIAC of weak electrolytes can reach very low values (weak salt-water interactions). Implementing a chemical-reaction approach according to Sec. 2.2 provokes a reduced number of free hydrated ions in the modelling, and consequently the calculated MIAC also decreases. 

Finally, it must be recalled that the transport properties of any material are strongly dependent on the molecular or ionic interactions, and that the dynamics of each entity are narrowly correlated with the neighboring particles. This is the main reason why the theoretical treatment of these processes often shows similarities with models used for thermodynamic properties. The most classical example is the treatment of dilute electrolyte solutions by the Debye-Hiickel equation for thermodynamics and by the Debye-Onsager equation for conductivity. 

We can see from Eig. 7.4, curve la, that this equation describes the experimental data in very dilute solutions of strong electrolytes (i.e., for 1 1 electrolytes approximately up to 10 M) for other electrolytes the concentration limit is even lower. It correctly conveys the functional dependence on the charge of the ions and the ionic strength of the solution (as well as the lack of dependence on individual properties of the ions) it can, moreover, be used to calculate the value of empirical constant h in Eq. (7.27). 

E. A. Guggenheim, Specific Thermodynamic Properties of Aqueous Solutions of Strong Electrolytes , Phil. Mag., 19, 588-643 (1935). 

In all other solutions the so called degree of dissociation, as determined from the measurement of some colligative property, merely indicates the magnitude of interionic forces, it cannot, however, be taken as a measure of the quantity of dissociated and undissociated molecules of the solute. A complete theory of strong electrolytes, at least of their diluted solutions, has been developed by Debye and Hiickel, this theory is the basis of modern electrochemistry. 

Weak electrolytes arc only partly dissociated into ions and their thermodynamic properties in solutions are influenced by electrostatic attractive forces to a much lesser degree, owing to a smaller ionic concentration, than in the ease of strong electrolytes. As hero it is a case of equilibrium between the dissociated and un dissociated parts of the electrolyte, the law of mass action can be applied to the reaction of dissociation. 

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