Freezing Point Lowering by Electrolytes

FREEZING POINT LOWERING BY ELECTROLYTES IN AQUEOUS SOLUTION... 

Freezing Point Lowering by Electrolytes in Aqueous Solution... 

The actual measurement of the freezing-point lowering caused by electrolytes has led us to choose the second postulate, that is, of the freedom of motion of the charged ions, as the nearer approximation to the true condition within the solution. 

When electrolyte solutions are involved, however, the osmotic effects such as freezing point lowering, osmotic pressure increase, and rise in boiling point are much greater than corresponds to the total electrolyte concentration. Accordingly, van t Hoff introduced the irrationality factor i (van t Hoff factor) by which the particular osmotic effect could be divided to yield a number which satisfied the equation of state (i is always greater than 1). The van t Hoff factor is purely empirical and does not account for the anomalous behavior of strong electrolyte solutions. 

Similarly, concepts of solvation must be employed in the measurement of equilibrium quantities to explain some anomalies, primarily the salting-out effect. Addition of an electrolyte to an aqueous solution of a non-electrolyte results in transfer of part of the water to the hydration sheath of the ion, decreasing the amount of free solvent, and the solubility of the nonelectrolyte decreases. This effect depends, however, on the electrolyte selected. In addition, the activity coefficient values (obtained, for example, by measuring the freezing point) can indicate the magnitude of hydration numbers. Exchange of the open structure of pure water for the more compact structure of the hydration sheath is the cause of lower compressibility of the electrolyte solution compared to pure water and of lower apparent volumes of the ions in solution in comparison with their effective volumes in the crystals. Again, this method yields the overall hydration number. 

Binary electrolytes, such as KC1, although completely ionized, even in the solid state, lower the freezing point less than 2 x 1.86D, even when as dilute as I0-3M. This was at first attributed to incomplete ioni/aiion but is now explained by the long range of electrostatic forces. Note that Mg++ and SO4 are less independent than K+ and Cl- AgNOa, unlike KC1, etc., is a weak salt, and undissociated molecules increase rapidly with concentration. The ions nearer to an ion of one sign arc those of opposite sign, therefore electric conductivity is less than the sum of ionic conductivities extrapolated to zero concentration. 

The thermodynamic treatment of systems in which at least one component is an electrolyte needs special comment. Such systems present the first case where we must choose between treating the system in terms of components or in terms of species. No decision can be based on thermodynamics alone. If we choose to work in terms of components, any effect of the presence of new species that are different from the components, would appear in the excess chemical potentials. No error would be involved, and the thermodynamic properties of the system expressed in terms of the excess chemical potentials and based on the components would be valid. It is only when we wish to explain the observed behavior of a system, to treat the system on the basis of some theoretical concept or, possibly, to obtain additional information concerning the molecular properties of the system, that we turn to the concept of species. For example, we can study the equilibrium between a dilute aqueous solution of sodium chloride and ice in terms of the components water and sodium chloride. However, we know that the observed effect of the lowering of the freezing point of water is approximately twice that expected for a nondissociable solute. This effect is explained in terms of the ionization. In any given case the choice of the species is dictated largely by our knowledge of the system obtained outside of the field of thermodynamics and, indeed, may be quite arbitrary. 

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. 

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