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  Home > JCE Print > Journal of Chemical Education > Issues > 2004  > November  >
In the Laboratory
The Effect of Ionic Strength on the Solubility of an Electrolyte
Joan D. Willey
Department of Chemistry and Marine Science Program, University of North Carolina at Wilmington, Wilmington, NC 28403-5932
Cover
November 2004
Vol. 81 No. 11
p. 1644
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.
Supplement
Background theory and instructions for students, procedures for the analysis of calcium by EDTA titration and by flame atomic absorption spectrophotometric analysis, and typical student results for these analyses are available.
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*  Citation
 Willey, Joan D. J. Chem. Educ. 2004 81 1644.
*  Keywords
 Aqueous Solution Chemistry; Demonstrations; Environmental Chemistry; Laboratory Instruction; Physical Chemistry; Quantitative Analysis; Water / Water Chemistry
*  History
 Created:
Last Updated:		 September 27, 2004
October 4, 2004
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