Prev. Article Next Article Table of Contents

Article

The Effect of Ionic Strength on the Solubility of an Electrolyte

Supporting Info

Your current credentials do not allow retrieval of the full text.

Purchase the full-text

PDF/HTML,
figures/images,
references and tables,
(where available)

Joan D. Willey

Department of Chemistry and Marine Science Program, University of North Carolina at Wilmington, Wilmington, NC 28403-5932

J. Chem. Educ., 2004, 81 (11), p 1644

DOI: 10.1021/ed081p1644

Publication Date (Web): November 1, 2004

Section:

History, Education, and Documentation

Abstract

The theory of activity versus concentration is important in industrial, environmental, and biochemistry. The increase in solubility of an electrolyte in a solution of a second electrolyte with no common ions compared with pure water is not an easy concept to grasp because it seems to be counterintuitive. The simple experiment described here illustrates this principle visually and dramatically. Students attempt to dissolve CaSO₄•2H₂O (gypsum) in pure water and in 0.25 M NaCl. The gypsum dissolves almost completely in the sodium chloride solution, but not in pure water. Students then measure the calcium concentrations in filtered aliquots of both solutions to quantify the solubility difference they observed. Students calculate mean activity coefficients using their measured concentrations and also from the Davies Equation, an extension of Debye–Hückel theory. The basic principle is there are ionic interactions between the solute ions and the solvent ions, which allow for more dissolution because only free ions enter into the expression for the solubility product equilibrium constant. From a simple mathematical point of view, in higher ionic strength solutions, activity coefficients for calcium and sulfate become smaller, and hence the concentrations must be larger to maintain a constant solubility product at equilibrium.