Diverse ion effect

Equation (7-4) indicates that the solubility product includes an activity-coefficient term, a term which has been assumed to be unity up to this time. The introduction to this chapter pointed out that errors arising from neglect of the effects of the activity coefficient are usually small when compared with several uncertainties or side reactions. The activity coefficient in Equation (7-4) depends on the kind and concentration of all electrolytes in solution, not merely those involved directly with the precipitate. The correction to solubility calculations that must be made to account for the activity-coefficient effect is known as the diverse ion effect. The appropriate background is discussed in Chapter 2, and Problems 2-1,2-2, and 2-3 are examples of the calculations. For 1 1 electrolytes in solution, activity coefficients can usually be assumed to be unity when concentrations are much less than 0.1 M. Common ion and diverse ion effects can be significant at the same time, for example, when a large excess of common ion is added in a precipitation. The diverse ion effect is one of the reasons that the haphazard addition of a large excess of precipitant should be avoided. 

Another important factor for coordination chemists is the general principle stated simply by Basolo solid salts separate from aqueous solution easiest for combinations of either small cation-small anion or large cation-large anion, preferably with systems having the same but opposite charges on the counter ions. Other factors governing solubility of salts are pH, interionic attraction, and the diverse ion effect, which will not be covered in detail here. 

Even though chemical reactions may go far to completion, the reactions never go in only one direction. In fact, reactions reach an equihbrium in which the rates of reactions in both directions are equal. In this chapter we review the equilibrium concept and the equilibrium constant and describe general approaches for calculations using equihbrium constants. We discuss the activity of ionic species along with the calculation of activity coefficients. These values are required for calculations using thermodynamic equihbrium constants, that is, for the diverse ion effect, described at the end of the chapter. They are also used in potentiometric calculations (Chapter 13). 

In our consideration of equilibrium constants thus far, we have assumed no diverse ion effect, that is, an ionic strength of zero and an activity coefficient of 1. Equilibrium constants should more exactly be expressed in terms of activities rather than concentrations. Consider the dissociation of AB. The thermodynamic equilibrium constant (i.e., the equilibrium constant extrapolated to the case of infinite dilution) Ktq is... 

Silver ion forms stepwise complexes with thiosulfate ion, 8203, with Kfi = 6.6 X 10 and = 4.4 X 10, Calculate the equilibrium concentrations of all silver species for 0.0100 MAgN03 in 1.00 M Na2S203. Neglect diverse ion effects. 

Figure 10.4 illustrates the increase in solubility of BuSOa in the presence of NaNOg due to the diverse ion effect. 

Calculate the solubihty of BaS04 in 0.0125 M BaCl2. Take into account the diverse ion effect. 

The diverse ion effect is more commonly called the salt effect. As a result of the salt effect, the numerical value of a K p based on molarities will vary depending on the ionic atmosphere. Most tabulated values of K p are based on activities rather than on molarities, thus avoiding the problem of the salt effect. 

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