The Effect of a Common Ion on Solubility

The presence of a common ion decreases the solubility of a slightly soluble ionic compound. As we saw in the case of acid-base systems, Le Chatelier s principle helps explain this effect. Let s examine the equilibrium condition for a saturated solution of lead(II) chromate  

After the addition, [Cr04 ] is higher, but [Pb ] is lower. In this case, [Pb ] represents the amount of PbCr04 dissolved thus, in effect, the solubility of PbCr04 has decreased. The same result is obtained if we dissolve PbCr04 in a Na2Cr04 solution. We also obtain this result by adding a soluble lead(II) salt, such 

Lead(il) chromate, a siightly soiubie salt, forms a saturated aqueous solution. 

When Na2Cr04 solution is added, the amount of PbCr04(s) increases. Thus, PbCr04 is less soluble in the presence of the common ion Cr04 . 

SAMPLE PROBLEM 19.7 Calculating the Effect of a Common Ion on Solubility 

SAMPLE PROBLEM 19.7 Calcularing ihe Effect of a Gxnmon Ion on Solubility 

Plan Addition of Ca , the common ion, should lower the solubility. We write the equation and ion-product expression and set up a reaction table, with [Ca ]i i, coming from Ca(N03)2 and S equal to [Ca Ifron, cacoHjj- To simplify the math, we assume that, because Ksp is low, S is so small relative to [Ca ]j i, that it can he neglected. Then we solve for S and check the assumption. 

---

Figure 19.10 The effect of a common ion on solubility. When a common ion is added to a saturated soiution of an ionic compound, the soiubiiity is iowered and more of the compound precipitates.

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. 

In Section 12.1, 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 solubihty product is an equilibrium constant precipitation 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 ][Cr] is, of course, equal to K p. Furthermore, simple stoichiometry tells us that [Ag ] = [Cl ]. But this equality does not hold in all situations. 

The Ion-Product Expression (2sp) and the Solubility-Product Constant (kfsp) 634 Calculations Involving the Solubility-Product Constant 635 Effect of a Common Ion on Solubility 637... 

Another aspect of the effect of electrolytes on the solubility of a salt is the concept of the solubility product for poorly soluble substances. The experimental consequences of this phenomenon are that if the concentration of a common ion is high, then the other ion becomes low in a saturated solution of the substance, that is, precipitation occurs. Conversely, the effect of foreign ions on the solubility of sparingly soluble salts is just the opposite, and the solubility increases. This is called the salt effect. 

We take advantage of the common ion effect to decrease the solubility of a precipitate in gravimetric analysis. For example, sulfate ion is determined by precipitating BaS04 with added barium chloride solution. Figure 10.3 illustrates the effect of excess barium ion on the solubility of BaS04. 

Of the possible substituting ions, COi ion is by far the most important followed by Na, S04 and Mg. The most common form of natural apatite in sedimentary rocks is francolite, a substituted form of carbonate fluorapatite deposited in marine systems. The substitution of col ior>s into the mineral lattice has a substantial effect on apatite solubility (Jahnke, 1984). More studies are required, however, before the effects of all substituting ions are imderstood and an accurate assessment of the solubility of complex, natural apatites can be made. 

If we proceed with part (b) as if it were an ordinary example of a common ion effect problem with Zn(OH)2 dissolving into a solution with a known concentration of OH- ion, we would get the wrong answer. We must take into account the effect of the formation of the complex ion, Zn(OH)42, on the solubility of Zn(OH)2. 

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. 

By far the most abundant phosphate mineral is apatite, which accounts for more than 95% of all P in the Earth s crust. The basic composition of apatite is listed in Table 14-2. Apatite exhibits a hexagonal crystal structure with long open channels parallel to the "c" axis. In its pure form, F, OH, or Cl occupy sites along this axis to form fluorapatite, hydroxyapatite, or chlorapatite, respectively. However, because of the "open" nature of the apatite crystal lattice, many substitutions are possible and "pure" forms of apatite as depicted by the general formula in Table 14-2 are essentially never foxmd. Of the possible substituting ions, carbonate ion is by far the most important followed by Na, SO , and Mg " ". The most common form of natural apatite is francolite, a highly substituted form of carbonate fluorapatite deposited in marine systems. The substitution of CO3 ions into the mineral lattice has a substantial effect on apatite solubility (Jahnke, 1984). More studies are required, however, before the effects of all substituting ions are understood and an accurate assessment of the solubility of complex, natural apatites can be made. 

Figure 5a indicates the effect of the CTAB concentration on the rate constants of the complexes of 38b and 38c. In the case of the water soluble 38b ligand, the rate increases with increasing CTAB concentration up to a saturation level. This type of saturation kinetics is usually interpreted to show the incorporation of a ligand-metal ion complex into a micellar phase from a bulk aqueous phase, and the catalytic activity of the complex is higher in the micellar phase than in the aqueous phase. In the case of lipophilic 38c, a very similar curve as in Fig. 4 is obtained. At a first glance, there appears to be a big difference between these two curves. However, they are rather common in micellar reactions and obey the same reaction mechanism 27). 

Effect of Cation. Of all alkali chlorides, only lithium chloride is sufficiently soluble in ethylenediamine to act as an electrolyte. On the other hand, all alkali iodides are soluble in ethylenediamine. Since the anion has a large effect on current efficiency, a common anion such as the iodide ion must be used to compare the effect of the various cations. The results of runs 6 and 7 show that the current efficiency was slightly higher and the percentage of octalin formed much greater when rubidium iodide was used instead of lithium iodide. The metallic cations Li and Rb give markedly better current efficiency than the organic cations (runs 6, 7, 8, and runs 5, 9, 10, 11). 

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. 

---