Solubility equilibria effects

Common ion effect The tube at the left contains a saturated solution of silver acetate (AgC2H302). Originally the tube at the right also contained a saturated solution of silver acetate. With the addition of a solution of silver nitrate (AgNOs), the solubility equilibrium of the silver acetate is shifted by the common ion Ag+ and additional silver acetate precipitates. 

An important result of the concepts discussed in this section and the preceding one is that precipitation and complexation reactions exert joint control over metal ion solubility and transport. Whereas precipitation can limit the dissolved concentration of a specific species (Me ), complexation reactions can allow the total dissolved concentration of that metal to be much higher. The balance between these two competing processes, taking into account kinetic and equilibrium effects, often determines how much metal is transported in solution between two sites. 

By convention, [HA(s)] = [B(s)] = 1. Eqs. (6.1) represent the precipitation equilibria of the uncharged species, and are characterized by the intrinsic solubility equilibrium constant, Sq. The zero subscript denotes the zero charge of the precipitating species. In a saturated solution, the effective (total) solubility S, at a particular pH is defined as the sum of the concentrations of all the compound species dissolved in the aqueous solution ... 

Equation 1 implies that solubility is independent of solvent type, and is only a function of the equilibrium temperature and characteristic properties of the solid phase. In real systems the effect of non-ideality in the liquid phase can significantly impact the solubility. This effect can be correlated using an activity coefficient (y) to account for the non-ideal liquid phase interactions between the dissolved solute and solvent molecules. Eq. 1. then becomes [7,8] ... 

We ve already discussed the common-ion effect in connection with the dissociation of weak acids and bases (Section 16.2). To see how a common ion affects the position of a solubility equilibrium, let s look again at the solubility of MgF2 ... 

In the preceding chapter we have been considering the equilibrium of two phases of the same substance. Some of the most important cases of equilibrium come, however, in binary systems, systems of two components, and we shall take them up in this chapter. Wo can best understand what is meant by this by some examples. The two components mean simply two substances, which may be atomic or molecular and which may mix with each other. For instance, they may be substances like sugar and wrater, one of which is soluble in the other. Then the study of phase equilibrium becomes the study of solubility, the limits of solubility, the effect of the solute on the vapor pressure, boiling point, melting point, etc., of the solvent. Or the components may be metals, like copper and zinc, for instance. Then we meet the study of alloys and the whole field of metallurgy. Of course, in metallurgy one often has to deal with alloys with more than two components—ternary alloys, for instance, with three components—but they arc considerably more complicated, and we shall not deal with them. 

Infer the effect of heat on the solubility equilibrium of water and cornstarch. 

Static, equilibration studies indicated that a molecular association forms at the styrene/water interface in the presence of mixed emulsifiers. Interfacial layers were also observed in spinning drop experiments between various oil phases and aqueous mixed emulsifier solutions. The formation of these interfacial layers as a function of time was found to be a non-equilibrium effect that depended primarily on the chemical structure of the oil phase. Oil phase water solubility had little effect. 

Again, the common-ion effect causes both this equilibrium and the solubility equilibrium to shift to the right an increase in solubility results. 

If we wish to calculate the solubility of barium sulfate in a system containing hydronium and acetate ions, we must take into account not only the solubility equilibrium but also the other three equilibria. We find, however, that using four equilibrium-constant expressions to calculate solubility is much more difficult and complex than the simple procedure illustrated in Examples 9-4, 9-5, and 9-6. To solve this type of problem, the systematic approach described in Section llA is helpful. We then use this approach to illustrate the effect of pH and complex fonna-tion on the solubility of typical analytical precipitates. In later chapters, we use this same systematic method for solution of problems involving multiple equilibria of several types. 

When calculating gas partial pressures it quickly becomes apparent that the principal temperature effect lies in the change of Kh rather than in the activity coefficients. In Equation 9 a constant 5C°p for the solubility equilibrium is assumed. While this is unlikely to result in appreciable errors over the temperature range 0-40°C, further from 25°C the variation of 6C°p has a significant effect on Kh, requiring a more complex form of Equation 9 (55). However, the necessary data are not always available, particularly for C°p of the aqueous solute which generally accounts for most of the change. 

Excluded from the total solubility or effective Henry s law are subsequent reactions, for example, the oxidation of dissolved SO2 into sulfuric acid (sulfate). Such oxidation increases the flux or phase transfer into water and thereby the phase partitioning, but cannot be described by equilibrium conditions (Chapter 4.1.4). 

The effect of nanoscale porosity on the test phase diagram is considered. Pore walls are shown to have more effect on the equilibrium vapor composition than on methane solubility. The effect of the critical point shift in nanopores is demonstrated. 

The swelhng method may not determine solubility parameters effectively if the rigid crosslinked samples tend to chip or split on swelling. The method can be used for polymers with a physical network that contains knots formed by polymer crystalhtes. Samples of unplasticized PVC do not swell in ester plasticizers, but if the same polymer is hot-compounded with dioctyl phthalate and cooled, it can swell to equilibrium as does the amorphous crosslinked elastomer. Note that such swelling of PVC depends on the thermal treatment history of samples. 

It is further obvious from these considerations that the equilibrium between a polymer homologous mixture and a solution of this mixture in a solvent of low molecular weight is a very intricate affair. The composition of the polymer homologous mixture in the solid, (which in most cases is a gel containing a certain amount of the solvent), is different from that in the solution. At the addition of further amounts of polymer substance, the equilibrium is disturbed, in contrast to the solubility equilibrium in micromolecular substances. This is the so-called Bodenk6rper rule, according to which the solubility of a polymer depends on the total amount of polymer present. This rule is only apparent, however, since no such effect exists in the solubility of a well-fractionated sample. 

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