Common ion effects

The concentration of a particular ion in an ionic reaction can be increased by the addition of a compound which produces that ion upon dissociation. The 

Example 7. Calculate the sulphide ion concentration in a 0.25M hydrochloric acid solution saturated with hydrogen sulphide. 

This concentration has been chosen since it is that at which the sulphides of certain heavy metals are precipitated. The total concentration of hydrogen sulphide may be assumed to be approximately the same as in aqueous solution, i.e. 0.1 M the [H + ] will be equal to that of the completely dissociated HC1, i.e. 0.25M, but the [S2 ] will be reduced below 1 x 10 14 (see Example 6). 

Thus by changing the acidity from 1.0 x 10 M (that present in saturated H2S water) to 0.25 M, the sulphide ion concentration is reduced from 1 x 10 14 to 1.6 x 10 21. 

Example 8. What effect has the addition of 0.1 mol of anhydrous sodium acetate to 1 L of 0.1 M acetic acid upon the degree of dissociation of the acid  

However, upon dissolution, if r r+, A yd/f (cation) will be much more negative than Ahyd f°(anion) for all values of 

along a series of related salts with increasing r., but with r r+, Aiattice ° will remain nearly constant while Ahyd f° becomes less negative. Hence, Aso H° (and thus AsoiG°) will become less negative (equation 7.57) and solubility will decrease. 

Such a series is exemplified by the alkali metal hexa-chloroplatinates. The hydrated sodium salt has a very high solubility, while the solubilities of K2[PtCl6], Rb2[PtCl6] and Cs2[PtCl6) are 2.30 x 10 2.44 x lO and 

Although the above, and similar, arguments are qualitative, they provide a helpful means of assessing the pattern in solubilities for series of ionic salts. We stress ionic because equations 7.55 and 7.56 assume an electrostatic model. Our discussions in Section 6.15 and earlier in this section indicated how partial covalent character in silver halides affects solubility trends. 

Let us now return to equation 7.47, and relate the observed solubility of a salt to the magnitude of the difference between Aiattice r° and A ydG (equation 7.49), and in particular to the sizes of the ions involved. 

One exception is the so-called common ion effect observed in electrolyte solutions. Here, the solubility product governs the solubility behavior (Equations 6.1 and 6.2)  

The presence of a second electrolyte sharing a common ion with the dominant species will reduce the solubility. The solubility product Ki is a constant, and any increase in the concentration of one or the other component reduces the concentration of the counterion. As a result, the apparent solubility of the salt 

So far, we have discussed aqueous solutions containing a single, dissolved ionic salt, MX. Now we consider the effect of adding a second salt which has one of its ions in common with the hrst salt. 

If a salt MX is added to an aqueous solution containing the solute MY (the ion M is common to both salts), the presence of the dissolved M + ions suppresses the dissolution of MX compared with that in pure water this is the common-ion effect. 

The origin of the common-ion effect is seen by applying Le Chatelier s principle. In equation 6.58, the presence of Cff in solution (from a soluble salt such as KCl) will suppress the dissolution of AgCl, i.e. additional CP ions will shift the equilibrium to the left-hand side. 

The effect is analogous to that of mixing a weak acid with the salt of that acid (e.g. acetic acid and sodium acetate) to form a buffer solution. 

Initial aqueous ion concentrations / mol dm Equilibrium concentrations / moldm  

In the preceding sections, we looked at solutions that contained either a weak acid, a weak base, or a salt of a weak acid or base. In the remaining sections of this chapter, we will look at the effect of adding another solute to a solution of a weak acid or base. The solutes we will look at are those that significantly affect acid or base ionization— that is, strong acids and bases, and salts that contain an ion that is produced in the acid or base ionization. These solutes affect the equihbrium through the common-ion effect. 

Le Chatelier s principle was applied in Section 15.7 to the problem of adding substances to an equilibrium mixture. 

The common-ion effect is the shift in an ionic equilibrium caused by the addition of a solute that provides an ion that takes part in the equilibrium. Consider a solution of acetic acid, HC2H3O2, in which you have the following add-ionization equilibrium  

Suppose you add HCl(o ) to this solution. What is the effect on the add-ionization equihbrium Because HCl(a ) is a strong acid, it provides H30 ion, which is present on the right side of the equation for acetic acid ionization. According to Le Chatelier s principle, the equilibrium composition should shift to the left.  

The degree of ionization of acetic acid is decreased by the addition of a strong add. This repression of the ionization of acetic acid by HCl(a ) is an example of the common-ion effect. 

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Addition of excess H ions to this solution will cause the equilibrium to move towards undissociated acid thereby decreasing the concentration of Ac . This effect is known as the common-ion effect and is of considerable practical importance. Thus, e.g. in the precipitation of metal ions as insoluble sulphides. 

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. 

Sa.lts Salting out metal chlorides from aqueous solutions by the common ion effect upon addition of HCl is utilized in many practical apphcations. Typical data for ferrous chloride [13478-10-9] FeCl2, potassium chloride [7447-40-7] KCl, and NaCl are shown in Table 9. The properties of the FeCl2-HCL-H2 0 system are important to the steel-pickling industry (see Metal SURFACE TREATMENTS Steel). Other metal chlorides that are salted out by the addition of hydrogen chloride to aqueous solutions include those of magnesium, strontium, and barium. 

If (A i[X ]/A 2[Y ]) is not much smaller than unity, then as the substitution reaction proceeds, the increase in [X ] will increase the denominator of Eq. (8-65), slowing the reaction and causing deviation from simple first-order kinetics. This mass-law or common-ion effect is characteristic of an S l process, although, as already seen, it is not a necessary condition. The common-ion effect (also called external return) occurs only with the common ion and must be distinguished from a general kinetic salt effect, which will operate with any ion. An example is provided by the hydrolysis of triphenylmethyl chloride (trityl chloride) the addition of 0.01 M NaCl decreased the rate by fourfold. The solvolysis rate of diphenylmethyl chloride in 80% aqueous acetone was decreased by LiCl but increased by LiBr. ° The 5 2 mechanism will also yield first-order kinetics in a solvolysis reaction, but it should not be susceptible to a common-ion rate inhibition. 

These substances accelerate the reaction, and their effectiveness increases in the order given. This suggestion was questioned by Pocker, who found that the effects of such added substances were not directly proportional to their concentrations and could easily be explained by macro effects on the solvent character. He also found that common-ion effects were small in the reaction, the effect of added 1-methylpyridinium bromide was negligible, and that there was no evidence for surface catalysis on the walls of the vessel. There is an exact parallel between the relative rates of the Finkelstein reactions... 

The predictions of the pH/potential diagram are generally fulfilled, but in very concentrated acid solutions, attack may diminish, owing to the relative insolubility of the relevant salt in the acid. Thus, lead nitrate, although soluble in water, has (owing to common ion effect) only slight solubility in concentrated nitric acid, and the corrosion rate is reduced. Similarly, lead chloride is less soluble in moderately concentrated hydrochloric acid than... 

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. 

Click Coached Problems for a self-study module on the common ion effect... 

Sodium bicarbonate, 112-113 Sodium carbonate, 61 Sodium chloride 44q common ion effect and, 439 electrolysis, 499 formation, 3 structure, 36 Sodium chlorine, 4 Sodium hydroxide, 61,84,441 Sodium hypochlorite, 369-370 Sodium stearate, 595 Sodium vapor lamps, 135 Solids... 

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. 

This is a quadratic equation in [H + ] and may be solved in the usual manner. It can, however, be simplified by introducing the following further approximations. In a mixture of a weak acid and its salt, the dissociation of the acid is repressed by the common ion effect, and [H + ] may be taken as negligibly small by... 

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. 

Column chromatography see Chromatography Columns in gas chromatography, 238 in liquid chromatography, 223 Combustion flames 784 Common ion effect 26 quantitative effects of, 35 Comparators permanent colour standards,... 

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. 

We can gain a quantitative understanding of the common-ion effect by considering how a change in the concentration of one of the ions affects the solubility product. Suppose we have a saturated solution of silver chloride in water ... 

The prediction of the numerical value of the common-ion effect is difficult. Because ions interact with one another strongly, simple equilibrium calculations are rarely valid the activities of ions differ markedly from their molarities. However, we can still get an idea of the size of the common-ion effect by solving the Ksp expression for the concentration of the ion other than the common ion. 

The common-ion effect is the reduction in solubility of a sparingly soluble salt by the addition of a soluble salt that has an ion in common with it. 

J 10 Describe the common-ion effect and assess its magnitude (Example 11.9). 

If the X formed during the reaction can decrease the rate, at least in some cases, it should be possible to add X from the outside and further decrease the rate in that way. This retardation of rate by addition of X is called common-ion effect or the mass law effect. Once again, addition of halide ions decreases the rate for diphenylmethyl but not for tert-butyl halides. 

We have seen how the polarity of the solvent influences the rates of Sn 1 and Sn2 reactions. The ionic strength of the medium has similar effects. In general, the addition of an external salt affects the rates of SnI and Sn2 reactions in the same way as an increase in solvent polarity, though this is not quantitative different salts have different effects. However, there are exceptions though the rates of SnI reactions are usually increased by the addition of salts (this is called the salt effect), addition of the leaving-group ion often decreases the rate (the common-ion effect, p. 395). 

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