THE COMMON-ION EFFECT

The common ion effect may best be illustrated by considering the solubility of M Xj, in the presence of mol/L of anion X. Then, 

In most cases that will be encountered, bS, so that this equation simplifies to 

Calculate the molar solubilities, S, of AgCl and Ag2Cr04 under the following conditions. 

The common ion effect may be considered to apply to any solution in which there is more than one source of the ions contained in the precipitate. The common ion effect is involved in such questions as, what concentration of reagent anion would be necessary to initiate precipitation, and what is the anion concentration when the precipitation is (virtually) complete  

At this concentration, the solution is saturated no precipitate will form. When the concentration is slightly higher than this, precipitation will be initiated. 

The common ion effect is simply an application of Le Chdtelier s principle. 

Our discussion of acid-base ionization and salt hydrolysis in Chapter 15 was limited to solutions containing a single solute. In this section, we will consider the acid-base properties of a solution with two dissolved solutes that contain the same ion (cation or anion), called the common ion. 

The presence of a common ion suppresses the ionization of a weak acid or a weak base. If sodium acetate and acetic acid are dissolved in the same solution, for example, they both dissociate and ionize to produce CHsCOO ions  

The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance. The common ion effect plays an important role in determining the pH of a solution and the solubility of a slightly soluble salt (to be discussed later in this chapter). Here we will study the common ion effect as it relates to the pH of a solution. Keep in mind that despite its distinctive name, the common ion effect is simply a special case of Le ChStelier s principle. 

Let us consider the pH of a solution containing a weak acid, HA, and a soluble salt of the weak acid, such as NaA. We start by writing 

Recall that a system at equilibrium will shift in response to being stressed and that stress can be applied in a variety of ways, including the addition of a reactant or a product [ W Section 15.5]. Consider a liter of solution containing 0.10 mole of acetic acid. Using the for acetic add 

Now consider what happens when we add 0.050 mole of sodium acetate (CH3COONa) to the solution. Sodium acetate dis.sociates completely in aqueous solution to give sodium ions and acetate ions  

by adding sodium acetate, we have increased the concoitration of acetate ion. Because acetate ion is a product in the ionization of acetic add, the addition of acetate ion causes the equilibrium to shift to the left The net result is a reduction in the piercent ionization of acetic add. 

Shifting the equilibrium to the left consumes not only some of the added acetate ion, but also some of the hydrogen ion. This causes the pH to change (in this case the pH increases). Sample Problem 17.1 shows how an equilibrium table can be used to calculate the pH of a solution of acetic acid after the addition of sodium acetate. 

Student Annotation By adding sodium acetate, we also add sodium ions to the solution. However, sodium ions do not interact with water or with any of the other species present [Mt Section 16.10]. 

If a slightly soluble salt solution is at equilibrium and we add a solution containing one of the ions involved in the equilibrium, the solubility of the slightly soluble salt decreases. For example, consider the PbS04 equilibrium  

Suppose we add a solution of Na2S04 to this equilibrium system. The additional sulfate ion will disrupt the equilibrium by Le Chatclier s principle and shift it to the left. This decreases the solubility. The same would be true if you tried to dissolve PbS04 in a solution of Na2S04 instead of pure water—the solubility would be less. This application of Le Chatelier s principle to equilibrium systems of a slightly soluble salt is the common-ion effect. 

The solubility of stannous fluoride, SnFe(s), in water at 20 C is 0.012 g/100 ml. What is the solubility of SnFjjs) in a 0.08 M NaF solution neglecting ionic strength effects  

Now let S = moles/liter of SnFj, which will dissolve. If 8 x 10 mole of F-/liter is originally present, 

We see from this example that the solubility of Snp2 is reduced from 7.7 

X 10 mole/liter to 2.8 x 10 mole/liter by the presence of 0.08 M F . A similar calculation shown in Example 6-4 for a solution containing lO mole NaF/liter shows essentially no decrease in solubility caused by the common ion. 

What is the solubility of SnF in a 10 mole/liter NaF solution Neglect ionic strength effects. 

In Chapter 16 we examined the equilibrium concentrations of ions in solutions containing a weak acid ora weak base. We now consider solutions that contain a weak acid, such as acetic acid (CH3COOH), and a soluble salt of that acid, such as sodium acetate (CH3COONa). Notice that these solutions contain two substances that share a common ion, CH3COO. It is instructive to view these solutions from the perspective of Le Chatelier s principle. (Section 15.7) 

Sodium acetate is a soluble ionic compound and therefore a strong electrolyte. 3 - (Section 4.1) Consequently, it dissociates completely in aqueous solution to form Na and CH3COO ions  

In contrast, CH3COOH is a weak electrolyte that ionizes only partially, represented by 

When we have sodium acetate and acetic acid in the same solution, the CH3COO from CH3COONa causes the equilibrium of Equation 17.1 to shift to the left, thereby decreasing the equilibrium concentration of IV(ciq)  

In other words, the presence of the added acetate ion causes the acetic acid to ionize less than it normally would. 

Addition of CH3COO shifts equilibrium concentrations, lowering [H+] 

So far, we have discussed solubilities of salts in water. Now, the solubilities of salts in another salt solution will be discussed. If AgCl salt is added to a NaCl solution, the solubility of AgCl in the solution will be smaller than its solubility in pure water because the common - ion Cl causes decrease in solubility in the solution. 

Mixing the soiutions of iead(IV) nitrate and sodium iodide produces a yeiiow iead(ll) iodide precipitate. 

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  

W e can add a couple of different things to the mixture that will affect it and demonstrate the common-ion effect. One of them would be sodium acetate, NaC.H.O.. Another would be hydrochloric acid. Think about this for a moment, keeping in mind Le Chatelier s principle. If sodium acetate is added to the mixture, it will dissociate into sodium ions and acetate ions. The increase in concentration ofthe acetate ions will drive the reaction to the left, which will further inhibit the dissociation of acetic acid. Adding hydrochloric acid will have the same effect because it will increase the concentration of protons, which will also drive the reaction to the left. Sodium acetate and hydrochloric acid have two features that allow them both to cause the common-ion effect to occur. First, they are both strong electrolytes, and second they each have an ion in common with the acetic acid equilibrium. These are the key ingredients that cause the common-ion effect. 

Having a conceptual understanding of the effect is a good starting point, but we still need to be able to understand the quantitative relationships between the different components in the equilibrium mixture. In this section, we will see how to deal with the common-ion effect in acid-base equilibrium problems. You will find that these problems are very similar to the weak acid problems earlier in the chapter. 

In this sample, we will see what happens when we make a solution that contains both acetic acid and sodium acetate (which was used in our previous example). Suppose you make a solution that contains 0.30 mol HCjHjOj and 0.30 mol NaC H O dissolved in 1.00 liter of solution. Let s compare the pH of this sample to the pH of a solution that only has 0.30 mol HCjHjOj. If the conceptual explanation in the last section is valid, we should predict that the pH of the solution with the sodium acetate will be higher (less acidic) than the acetic acid-only solution. To begin, let s set up the equation for the equilibrium reaction  

There are a few things that you need to pay attention to as you set this problem up. First, the sodium need not be written in the equation because it doesn t do anything (it s a spectator). Second, unlike the previous weak acid problems, these problems aren t starting out with no products. The common ion in this reaction is the acetate ion, which is a product of the acetic acid dissociation. When you set up your chart, you need to include all amounts of all substances present at the start of the reaction. We re going to omit water in our chart because it is not part of the equilibrium expression. 

We can add a couple of different things to the mixture that will affect it and demonstrate the common-ion effect. One of them would be sodium acetate, NaC2H302. Another would be hydrochloric acid. Think about this for a moment, keeping in mind Le Chatelier s Principle. If sodium acetate is added to the mixture, it will dissociate into sodium ions and acetate ions. 

Increases in the concentrations of calciunj, bicarbonate or carbonate from sources other than the dissolution of calcite may supersaturate a water with respect to calcite, causing it to precipitate. In the Floridan carbonate aquifer, groundwaters attain calcite saturation by contact with limestone, but then may become supersaturated with respect to calcite because of gypsum dissolution (Back and Han-shaw 1970 Wicks and Herman 1996). Calcium from the gypsum, which is far more soluble than calcite, drives the common ion effect reaction 

CaS04 2H20(gypsum) + 2HCO3- CaC03 + SO + 2H2O + CO2 (6.27) 

If the common ion is calcium, then the reaction is described by a vertical arrow pointing upward in Fig. 6.7(d) (line A-B). With further dissolution of gypsum and Ca increases, groundwater composition then moves upward along the calcite saturation curve (line B-C) in Fig. 6.7(d) to higher CO2 pressures. Point C will correspond to saturation with respect to gypsum, if the mineral is not completely leached from the rock. 

Unlike the aluminosilicates and most other minerals, the carbonates have an exothermic heat of dissolution, which means that their solubilities decrease with increasing temperature. For example, for calcite decreases from 10 at 0°C to 10 at 30°C. The effect of temperature on the solubility products of aragonite, calcite, and ordered dolomite is plotted in Fig. 6.8. The figure shows that between 0 and 90°C solubilities of the carbonates decrease by about 6-fold for aragonite and calcite and 14-fold for dolomite. This decreasing solubility with temperature is magnified by the fact that the solubility of CO2 gas also declines with temperature. Kqq, decreases from 10 to 10 between 0 and 30°C. 

Mixing of two waters can lead to either carbonate mineral supersaturation or undersaturation of their mixture. The effect can be simply understood for calcite from a plot of Ca + versus C02(aq) (which, by convention, equals H2CO3). On such a plot, shown in Fig. 6.9, saturation with respect to calcite is defined by a curve, whereas mixtures of two waters lie on a straight line connecting their compositions. If we mix saturated waters of compositions A and B in Fig. 6.9, regardless of their 

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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. 

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. 

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

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. 

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). 

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). 

The concentrations of ions in Example reveal an important feature of aqueous equilibria. Notice in the calculations that the amount of dissolved Cd ions decreased by a factor of 10 upon addition of excess OH. This reduction in concentration is an example of the common-ion effect. 

Figure 18-9 is a molecular illustration of the common-ion effect, and Example shows a quantitative application. 

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 ". ... 

The common-ion effect is quite general. The chemistry of buffer solutions is another important application of this principle. A buffer solution relies on the common-ion effect to suppress the concentration of hydronium ions and maintain a steady pH ... 

One of the characteristics of complexation equilibria is that, in the presence of excess concentration of a complexing ligand, formation of a complex often reduces the concentration of a free metal cation essentially to zero. This is another application of the common-ion effect discussed in the previous section. Example treats a situation of this sort. 

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. 

The common ion effect alters the amount of solid that will dissociate in solution. The addition of solid silver chromate to an aqueous solution of potassium chromate will affect the silver chromate s solubility because —... 

A solids are always subject to the common ion effect in solution... 

The removal of dissolved inorganic compounds is usually accomplished by precipitation. This step is based on the common ion effect. When a salt dissolves in water, it forms two ions. The salt will continue to dissolve in water until the product... 

All efforts of trying to extract the sodium salt into an organic solvent were unsuccessful, and, thus, a salting out procedure using the common ion effect was used. The product at this stage is in the aqueous phase and addition of 15 g of sodium chloride /100 mL of the reaction mixture, completely precipitates all the coupled sodium salt which is filtered. The filtered crude product is dewatered by washing with isopropyl alcohol or drying under vacuum. The mother liquors... 

Thus, by an application of the common ion effect, the solubility of barium sulfate has been reduced to produce 104-fold less free barium ion in solution, thus further reducing the risk of barium toxicity. 

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