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An Analytical Chemistry Experiment
Solubility equilibria are usually introduced in the general chemistry course with limited mention of the nonideal behavior of real solutions. Nonideal behavior is introduced in the quantitative analytical chemistry course in the context of gravimetry. Often this concept is left as an academic exercise with no experimental application. The determination of a solubility-product constant as reported here differs from other experiments (1-3) in that nonideality is taken into account through Debye-Huckel theory by making few simplifying assumptions.
The measurement of a thermodynamic constant is a valuable and satisfying experience for students who may otherwise find lecture concepts of nonideality intimidating or unmanageable. One recent article reports the experimental determination of the thermodynamic acid-dissociation constant (4). The project reported here studies the effect of ionic strength on the solubility of lead(II) iodide and models the effect with Debye-Huckel theory. From measured solubility in solutions of varying ionic strength, the thermodynamic solubility product, Ksp, is calculated using activity coefficients derived from the extended DebyeHuckel law and the solubility at infinite dilution assuming the Debye-Huckel limiting law is an acceptable model. The apparent solubility product, *Sp, is calculated assuming solution ideality and then compared with the calculated Ks .
The equilibrium mass-action equation for the sparingly soluble lead(II) iodide in aqueous solution is written The thermodynamic solubility-product expression, written in terms of activity, a, is or, in terms of molar concentration, where fpb2 and fI are the activity coefficients of the lead and iodide ions.
Solution ideality can often be assumed in calculations involving solubility equilibria. In the ideal case, f is assumed to be unity so that the solubility, s (=[Pb"]), of PbI2 is calculated as Kp/4Theactivitycoefficientthereforesolubility) is a function of ionic strength, gt; ionic charge, Z; and effective ionic diameter, R. Effective ionic diameters are tabulated in most chemical reference books and quantitative analytical chemistry texts (5, 6). The value of f for a given ion, i, can be calculated with the extended DebyeHickel expression, By nature of the electrostatic behavior of ions in aqueous solution, larger effective ionic diameters and higher ionic charges lead to larger deviations from ideality (7, 8). In the more rigorous solution, solubility is calculated as If solution ideality is not assumed, it is possible to...