Solubility common-ion effect and

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. 

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. 

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. 

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. 

EXAMPLE 3 H2S will not precipitate FeS from a strongly acid (HC1) solution of Fe2+. The large [H+] furnished by the hydrochloric acid stalls the ionization of H2S (common-ion effect) and reduces the [S2-] to so low a value that the solubility product of FeS is not reached. 

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... 

However, factors such as the solubility product (Ksp) of the salt, common-ion effects, and hygroscopi-city may disfavor the salts of first choice using the above criterion, e.g., hydrochloride or sodium salts. The formation of hydrochloride salts does not always enhance solubility above that of the free base. The lower solubility of a hydrochloride salt in dilute HCl, relative to that of the free base, is attributed to the common ion effect of the chloride ion on the solubility product equilibrium of the salts. The common-ion effect suppresses the solubility product equilibrium. This is particularly relevant to the HCl salts of drugs administered orally, resulting in contact with... 

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... 

If there is an excess of one ion over the other, the concentration of the other is suppressed (common ion effect), and die solubility of the precipitate is decreased. We can still calculate the concentration from the solubihty product. 

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

This would be a messy problem to solve exactly, but fortunately it is possible to simplify matters. Even without the common-ion effect, the solubility of Cap2 is very small (2.1 X 10 M). Thus, we assume that the 0.010 M concentration of Ca from Ca(N03)2 is very much greater than the small additional concentration resulting from the solubility of Cap2 that is, X is much smaller than 0.010 M, and 0.010 + X — 0.010. We then have... 

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. 

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