Ion effects on solubility

So far we have only paid attention to ionic species and their solubility in pure water. In this section we shall look at how solubility is influence by the presence of ions that one way of another also participates in the reaction of dissolution. This is something that we will look into in the following example  

We wish to determine the solubility of silver ehromate Ag2Cr04 in a 0.100 M aqueous solution of AgNOs. Silver ehromate has a solubility produet of Ksp = 9.0 lO before the solid silver ehromate is dissolved there is already Ag and NO3 ions present in the solution. As NO3 does not partieipate in the dissolution reaetion for silver ehromate, we may ignore this ion. The initial eoneentration of Ag of 0.100 M has nevertheless an importanee. We have the following initial eoneentrations  

Ksp = [Ag ] [Cr04T = 9.0 lO M We assume as often thatx moles/L of Ag2Cr04(s) should be dissolved in order to reaeh equilibrium  

There may thereby by dissolved 9.0 10 ° mol/L Ag2Cr04(s) in 0.100 M AgNOs. It is easy to realise that had we had a pure water solution, the solubility would have been found by the following equation  

there may be dissolved 1.3 lO moles/L of Ag2Cr04(s) in pure water. It may thereby been seen that by comparison of the aqueous solution with the 0.100 M AgNOs solution that there may be dissolved far more Ag2Cr04(s) in pure water then in the solution containing Ag ions already present. One may in an informal way say that the silver ions already present in the solution hinder the dissolution of silver chromate. 

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Serajuddin, A. T. M., P.-C. Sheen, and M. A. Augustine. 1987. Common ion effect on solubility and dissolution rate of the sodium salt of an organic acildPharm. PharmacoB9 587-591. 

Common ion effect on solubility Overnight equilibration at 25°C in suitable media and analysis by UV-VIS or HPLC Compare solubility in demineralized water with 1.2%wAr saline for salts and parent... 

AgCIfe) -> Ag+(aq) + C (aq) some NaCl is added to the solution Shift to left due to increase in Cl- concentration. This is known as the common ion effect on solubility. 

The equilibrium constant for the equilibrium between a shghtly soluble ionic solid and its ions in solution is called the solubility product constant, K p. Its value can be determined from the solubility of the solid. Conversely, when the solubility product constant is known, the solubility of the solid can be calculated. The solubility is decreased by the addition of a soluble salt that supplies a common ion. Qualitatively, this can be seen to follow from Le Chatelier s principle. Quantitatively, the common-ion effect on solubility can be obtained from the solubility product constant. 

For the estimation of medium effects on solubility equilibria in mixtures of electrolytes involving ions of charge less than three, one may neglect 6 and j/ terms in Eq. (6.33). 

Complex ions of the type AgX2 and AgX2- have been described for chloride, bromide and iodide and their relative stability was found to be I- > Br > Cl-.360 The effect of this complex ion formation on solubility can be readily assessed from the observation that in 1M HC1, AgCl becomes —100 times more soluble than in pure water. 

Dissolution of Al(OH)3 in excess base is just a special case of the effect of complex-ion formation on solubility Al(OH)3 dissolves because excess OH - ions convert it to the soluble complex ion Al(OH)4- (aluminate ion). The effect of pH on the solubility of Al(OH)3 is shown in Figure 16.16. 

The smallness of the ion effect on the water spectra is contradictry to the large effects by ions on the solubility of organic molecules in water. The apparent paradox can be easily understood by the simple water model. At room T the content of orientation defects is not very sensitive to the expansion of H-bond systems12,49 ... 

Negative values ofN —N0, the electrolyte effect on the association numbers of water, are called the structure-breaker effect. One can speak of negative hydration31. The estimation of the hydration numbers by spectroscopic or solubility methods gives only an approximation of the sum effect. The spectra of the H-bond bands show in second approximation distinct differences between the ion effects on the H-bonds7 ). — The partial molar volume Vx of water in electrolyte solutions is negative in all solutions but the series of -values corresponds to the Hofmeister ion series too. The negative V1 volume indicates an electrostriction effect around the ions. 

The solubility product of the Hg2Cl2 (calomel) is very low (Ksp = 1.3 x 10 l7). The potential of this electrode is again determined by the concentration of the chloride ion in the inner compartment. When a saturated solution of KC1 is used its potential against the SHE is n = +241 mV. Use of a saturated KC1 solution hides a certain danger the higher temperature sensitivity, which is due to the temperature effect on solubility. 

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 ... 

FIGURE 16.14 An illustration of the effect of complex ion formation on solubility. Each test tube contains 2.0 g AgBr, but the one on the left also contains dissolved thiosulfate ion, which forms a complex ion with Ag. Almost none of the white solid AgBr has dissolved in pure water, but all of it has dissolved in the solution containing thiosulfate. 

In Section 16.2 we discussed the effect of a common ion on acid and base ionizations. Here we will examine the relationship between the common ion effect and solubility. 

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