Acid-base equilibria containing common ions

THE COMMON ION EFFECT (SECTION 17.1) In this chapter, we have considered several types of important equilibria that occur in aqueous solution. Our primary emphasis has been on acid-base equilibria in solutions containing two or more solutes and on solubility equilibria. The dissociation of a weak acid or weak base is repressed by the presence of a strong electrolyte that provides an ion common to the equilibrium (the common-ion effect). 

Recall, from Chapter 8, that a buffer consists of a weak acid/conjugate base mixture or a weak base/conjugate acid mixture. One buffer that you examined previously contains acetic acid and sodium acetate. The common-ion effect applies to this buffer. The equilibrium of the acetic acid is affected by the common acetate ion from sodium acetate. 

The solubility product is the equilibrium constant for the dissolution of a solid salt into its constituent ions in aqueous solution. The common ion effect is the observation that, if one of the ions of that salt is already present in the solution, the solubility of a salt is decreased. Sometimes, we can selectively precipitate one ion from a solution containing other ions by adding a suitable counterion. At high concentration of ligand, a precipitated metal ion may redissolve by forming soluble complex ions. In a metal-ion complex, the metal is a Lewis acid (electron pair acceptor) and the ligand is a Lewis base (electron pair donor). 

You will see this topic appear twice, once in this chapter and once the next chapter. For now, you will see how this phenomenon affects acid-base equilibria. In the next chapter, you will see its effects on solubility equilibria. The common-ion effect is not too different from what its name suggests. If you have an equilibrium system and add a solute to it that contains one of the ions in the equilibrium, it will cause the equilibrium to shift. That is the common-ion effect (common because the solute has an ion in common with the equilibrium system). From a conceptual standpoint, this can be addressed using Le Chatelier s Principle. For example, consider our favorite equilibrium system below ... 

The most common procedure is to carry out the measurements in the form of a titration. Most commonly a solution containing metal-ion, ligand and acid is titrated with base. Occasionally, when the rate of attaining equilibrium in the system is slow, a "batch" method is adopted individual solutions of appropriate concentrations are prepared, sealed, placed in a constant-temperature bath and allowed to reach equilibrium, at which time the final pH measurements are made. 

However, the case in which the solubility of a solid can be calculated from the known analytical concentration of added components and from the solubility product alone is very seldom encountered. Ions that have dissolved from a crystalline lattice frequently undergo chemical reactions in solution, and therefore other equilibria in addition to the solubility product have to be considered. The reaction of the salt cation or anion with water to undergo acid-base reactions is very common. Furthermore, complex formation of salt cation and salt anion with each other and with one of the constituents of the solution has to be considered. For example, the solubility of FeS(s) in a sulfide-containing aqueous solution depends on, in addition to the solubility equilibrium, acid-base equilibria of the cation (e.g., Fe + H2O = FeOH + H ) and of the anion (e.g., S + HjO = HS + OH, and HS" + H2O = H2S + OH ), as well as on equilibria describing complex formation (e.g., formation of FeHS" or FeSi ). 

It is obvious that the expression enclosed in the brackets by the author of the present book is nothing but the primary medium effect of O2- expressed via the difference in the values of the equilibrium constants of equation (1.3.6) for the media compared the molten equimolar KCl-NaCl mixture, which was chosen as a reference melt, and for which pKHa/H20 was found to be 14 at 700 °C, and the melt studied. As to the physical sense of the common acidity function Cl, this is equal to the pO of the solution in the molten equimolar KCl-NaCl mixture, whose acidic properties (oxide ion activity) are similar to those of the solution studied. Moreover, from equation (1.3.7) it follows that solutions in different melts possess the same acidic properties (f ) if they are in equilibrium with the atmosphere containing HC1 and H20 and Phc/Ph2o — constant. This explanation confirms that the f function is similar to the Hammett function. Therefore, Cl values measured for standard solutions of strong bases in molten salts allow the prediction of the equilibrium constants on the background of other ionic solvents from the known shift of the acidity scales or the f value for the standard solution of a strong Lux base in the solvent in question. According to the assumption made in Refs. [169, 170] this value may be obtained if we know the equilibrium constant of the acid-base reaction (1.3.6) in the solvent studied. 

In Chapter 16 we examined the equilibrium concentrations of ions in solutions containing a weak acid ora weak base. We now consider solutions that contain a weak acid, such as acetic acid (CH3COOH), and a soluble salt of that acid, such as sodium acetate (CH3COONa). Notice that these solutions contain two substances that share a common ion, CH3COO. It is instructive to view these solutions from the perspective of Le Chatelier s principle. (Section 15.7)... 

The dissociation of the weak acid HC2H3O2 decreases when we add the strong electrolyte NaC2H302, which has an ion in common with it. We can generalize this observation, which we call the common-ion effect. The extent of ionization of a weak electrolyte is decreased by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte. Sample Exercises 17.1 and 17.2 illustrate how equilibrium concentrations may be calculated when a solution contains a mixture of a weak electrolyte and a strong electrolyte that have a common ion. The procedures are similar to those encountered for weak acids and weak bases in Chapter 16. 

Solutions of Acids or Bases Containing a Common Ion Equilibrium Calculations... 

The questions answered in Chapter 16 were mostly of the type, "What is the pH of 0.10 M CH3COOH, of 0.10 M NH3, of 0.10 M H3PO4, of 0.10 M NH4CI " In each of these cases, we think of dissolving a single substance in aqueous solution and determining the concentrations of the species present at equilibrium. In most situations in this chapter, a solution of a weak acid or weak base initially contains a second source of one of the ions produced in the ionization of the acid or base. The added ions are said to be common to the weak acid or weak base. The presence of a common ion can have some important consequences. 

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