Conductimetric determination of thermodynamic pairing constants for symmetrical electrolytes.

Earlier theories of electrolytic conductance are reviewed; all of these, with the exception of the Arrhenius-Ostwald theory, are based on physical models. Their theory failed to describe the conductance of strong electrolytes because it did not include the effects (then unsuspected) of long-range forces on mobility. Thermodynamic derivations are independent of model; applied to the postulated equilibrium A(+) + B(-) right arrow over left arrow A(+)B(-) between free ions and nonconducting paired ions, the thermodynamic pairing constant K(a) equals a(p)/(a+/-)(2), and DeltaG, the difference in free energy between paired ions (activity = a(p)) and free ions (activity = a(+/-)), equals (-RT ln K(a)). Converting to the molarity scale, K(a) = (1000 rho/M)[1 - gamma)/cy(2)(y(+/-))(2)]. Here rho is the density of the solvent of molecular weight M, c is stoichiometric concentration of electrolyte (mol/liter), gamma is the fraction of solute present as unpaired ions, and y(+/-) is their activity coefficient. The corresponding conductance function Lambda = Lambda(c;Lambda(0),R, big up tri, openG)involves three parameters: Lambda(0), the limiting equivalent conductance; R, the sum of the radii of the cospheres of the ions; and DeltaG. Conductance data for cesium bromide and for lithium chloride in water/dioxane mixtures and for the alkali halides in water are analyzed to determine these parameters. Correlations between the values found for R and DeltaG and properties characteristic of salt and solvent are then discussed.