Solubility and the Common-Ion Effect

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. 

As we have noted, the solubility product is an equilibrium constant predpitalion of an ionic compound from solution occurs whenever the ion product exceeds for that substance. In a saturated solution of AgCl, for example, the ion product [Ag ][Cl ] is, of course, equal to K p. Furthermore, simple stoichiometry tells us that [Ag ] = [CU]. But this equality does not hold in all situations. 

Suppose we study a solution containing two dissolved substances that share a common ion, say, AgCl and AgNOs. In addition to the dissociation of AgCl, the following process also contributes to the total concentration of the common silver ions in solution  

If AgNOs is added to a saturated AgCl solution, the increase in [Ag ] will make the ion product greater than the solubility product  

To reestablish equilibrium, some AgCl will precipitate out of the solution, as Le ChStelier s principle would predict, until the ion product is once again equal to Kgp. The effect of adding a common ion, then, is a decrease in the solubility of the salt (AgCl) in solution. Note that in this case [Ag ] is no longer equal to [CU] at equilibrium rather, [Ag ] [CU]. 

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Another crystallization technique is used when the isolation of a highly water-soluble compound in its salt form is required from aqueous reaction mixtures. This technique takes advantage of the common-ion effect and is based on the le Chatelier s principle, which states that, if, to a system in equilibrium, a stress is applied, the system will react so as to relieve the stress. Thus, in aqueous solutions, the solubility of the compound in salt form can be reduced by adding large amoimts of a common ion which is more soluble than the salt of the compoimd. 

The effects of various factors such as pH, the common ion effect, and temperature on solubility will have a greater impact on formulation development for insoluble compounds than for soluble ones. The general solubility theory has been extensively discussed in the literature (James, 1986 Grant and Higuchi, 1990). To afford better understanding of the solubility behavior of insoluble compounds, the pertinent solubility theory and its practical implications will be reviewed here. 

Solubility of Sparingly Soluble Salts, the "Common Ion" Effect and Le Chatelier s Principle... 

Parentheses denote activity and brackets denote concentration of the species. The concentration of the Al(OH)3 species represents approximately the lowest possible solubility point of the mineral and it is the product of two constants (K -K ). Thus, its magnitude is not in any way related to pH. Mineral solubility increases as pH increases above the solution pH of zero net charge because of increasing complexa-tion effects, and mineral solubility also increases at pH values below the solution pH of zero net charge because of diminishing common-ion effects (Fig. 2A). All minerals are subject to the common-ion effect and many minerals are subject to the complexation or ion-pairing effect (Fig. 2B). 

Because there were already sulfate ions in solution and more sulfate ions were added, sulfate ion is called the common ion. Adding sodium sulfate to the solution increases the concentration of sulfate ion in solution driving the reverse reaction. This is called the common ion effect and more of the solid calcium sulfate will be made. If the solid is being formed that means that it is not dissolving and the solubility has decreased. 

Figure 7-2 represents data for solubility of silver chloride plotted to illustrate the minimum solubility, the common ion effect and intrinsic solubility S°, and the formation of the chloro complex. As Figure 7-25 indicates, at low concentrations of excess chloride an essentially linear relation is obtained between solubility and the reciprocal of the product of chloride ion concentration and the square of the activity coefficient. The zero intercept corresponds to S° (in this case an intrinsic solubility... 

We investigate some of the factors that affect solubility, including the common-ion effect and the effect of acids. 

The addition of an electrolyte to a saturated solution of a sparingly soluble salt with a common ion depresses the solubility of the latter (the common ion effect) and leads to its precipitation. For example, the solubility product of AgCl at 25 °C is 1.56 X 10 ° (Ka = Kc at this dilution), i.e. 

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. 

The solubility of PbCU is affected by other chlorides, falling to a minimnm (the common-ion effect) and then rising again with c(CH (formation of complexes... 

Comparison of the common-ion effect and the salt effect on the molar solubility of Ag2Cr04... 

Sodium sulphate crystallises out in hydrated form (common ion effect) and is filtered off on concentration, sodium dichromate is obtained. For analytical purposes, the potassium salt. K2Cr20-. is preferred potassium chloride is added and the less soluble potassium dichromate obtained. 

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. 

The solubility of the precipitates encountered in quantitative analysis increases with rise of temperature. With some substances the influence of temperature is small, but with others it is quite appreciable. Thus the solubility of silver chloride at 10 and 100 °C is 1.72 and 21.1mgL 1 respectively, whilst that of barium sulphate at these two temperatures is 2.2 and 3.9 mg L 1 respectively. In many instances, the common ion effect reduces the solubility to so.small a value that the temperature effect, which is otherwise appreciable, becomes very small. Wherever possible it is advantageous to filter while the solution is hot the rate of filtration is increased, as is also the solubility of foreign substances, thus rendering their removal from the precipitate more complete. The double phosphates of ammonium with magnesium, manganese or zinc, as well as lead sulphate and silver chloride, are usually filtered at the laboratory temperature to avoid solubility losses. 

Calcium oxalate monohydrate has a solubility of 0.0067 g and 0.0140 g L 1 at 25° and 95 °C respectively. The solubility is less in neutral solutions containing moderate concentrations of ammonium oxalate owing to the common-ion effect (Section 2.7) hence a dilute solution of ammonium oxalate is employed as the wash liquid in the gravimetric determination. 

We can use Le Chatelier s principle as a guide. This principle tells us that, if we add a second salt or an acid that supplies one of the same ions—a common ion —to a saturated solution of a salt, then the equilibrium will tend to adjust by decreasing the concentration of the added ions (Fig. 11.15). That is, the solubility of the original salt is decreased, and it precipitates. We can conclude that the addition of excess OH- ions to the water supply should precipitate more of the heavy metal ions as their hydroxides. In other words, the addition of OH ions reduces the solubility of the heavy metal hydroxide. The decrease in solubility caused by the addition of a common ion is called the common-ion effect. 

Both solubilities are low, as we would expect for a salt with a small value of. S sp. Notice that PbCl2 is about 350 times less soluble in the NaCl solution. This makes sense in terms of the common-ion effect. The excess chloride ion suppresses the solubility of Pb by Le Chatelier s principle. The actual concentration of lead in seawater is much less than 4.0 X 10 M. This is because other lead salts are much less soluble than lead(II) chloride. The ocean contains carbonate, for example, and. STsp for lead(II) carbonate is quite small, 7.4 X lO ". ... 

In discovery the most common base salt is the HCl salt. It is common in biology to use buffers containing chloride and the chloride content of the gastric contents is about 0.15 M. It often happens that the common ion effect of chloride suppresses the solubility of an HCl salt. This is a 100% solvable problem in pharmaceutical sciences. 

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