The Solubility of Ionic Compounds

As we have just discussed, when an ionic compound dissolves in water, the resulting solution contains not the intact ionic compound itself, but its component ions dissolved in water. However, not all ionic compounds dissolve in water. If we add AgCl to water, for example, it remains solid and appears as a white powder at the bottom of the water. 

In general, a compound is termed soluble if it dissolves in water and insoluble if it does not. However, these classifications are a bit of an oversimplification. (In reality, solubility is a continuum and even insoluble compounds dissolve to some extent, though usually orders of magnitude less than soluble compounds.) For example, silver nitrate is soluble. If we mix solid AgN03 with water, it dissolves and forms a strong electrolyte solution. Silver chloride, on the other hand, is almost completely insoluble. If we mix solid AgCl with water, virtually all of it remains as a solid within the liquid water. 

A AgCl does not dissolve in water it remains as a white powder at the bottom of the beaker. 

Whether a particular compound is soluble or insoluble depends on several factors. In Section 12.3, we will examine more closely the energy changes associated with solution formation. For now, we can follow a set of empirical rules that chemists have inferred from observations on many ionic compounds. These solubility rules are summarized in Table 4.1. 

TABLE 4.1 Solubility Rules for Ionic Compounds in Water 

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As we noted in Chapter 4, the solubility of ionic compounds in water varies tremendously from one solid to another. The extent to which solution occurs depends on a balance between two forces, both electrical in nature ... 

Hard water is the name given to water supplies that contain significant concentrations of Mg2+ and Ca2+ ions. Check on the solubility of ionic compounds formed with these ions and predict what problems they may cause. 

In this chapter, you will continue your study of acid-base reactions. You will find out how ions in aqueous solution can act as acids or bases. Then, by applying equilibrium concepts to ions in solution, you will be able to predict the solubility of ionic compounds in water and the formation of a precipitate. 

In this section, you learned why solutions of different salts have different pH values. You learned how to analyze the composition of a salt to predict whether the salt forms an acidic, basic, or neutral solution. Finally, you learned how to apply your understanding of the properties of salts to calculate the pH at the equivalence point of a titration. You used the pH to determine a suitable indicator for the titration. In section 9.2, you will further investigate the equilibria of solutions and learn how to predict the solubility of ionic compounds in solution. 

Typically, buffers in the region of pH 7-9 have been used in MEEKC. At these pH values the buffers generate a high electroosmotic flow (EOF). Extreme values of pH have been used in MEECK specifically to suppress solute ionization. For example, a pH of 1.2 of the buffer has been used to prevent the ionization of acids (30,31). To eliminate the ionization of basic compounds, a buffer at pH 12 has been used. These pH values were used in MEEKC to measure the solubility of ionic compounds (30). High-pH carbonate buffers (31) were applied in place of the standard borate or phosphate buffers. 

Many different types of reversible reactions exist in chemistry, and for each of these an equilibrium constant can be defined. The basic principles of this chapter apply to all equilibrium constants. The different types of equilibrium are generally denoted using an appropriate subscript. The equilibrium constant for general solution reactions is signified as or K, where the c indicates equilibrium concentrations are used in the law of mass action. When reactions involve gases, partial pressures are often used instead of concentrations, and the equilibrium constant is reported as (p indicates that the constant is based on partial pressures). and are used for equilibria associated with acids and bases, respectively. The equilibrium of water with the hydrogen and hydroxide ions is expressed as K. The equilibrium constant used with the solubility of ionic compounds is K p. Several of these different K expres-... 

The roles of hydration enthalpies and entropies in determining the solubilities of ionic compounds... 

Equilibria govern diverse phenomena from the folding of proteins to the action of acid rain on minerals to the aqueous reactions used in analytical chemistry. This chapter introduces equilibria for the solubility of ionic compounds, complex formation, and acid-base reactions. Chemical equilibrium provides a foundation not only for chemical analysis, but also for other subjects such as biochemistry, geology, and oceanography. 

For a vary thorough discusson of enthalpy, entropy, and the solubility of ionic compounds, see Johnson. D. A. Some Thermodynamic Aspects of Inorganic Chemistry Cambridge University London. 1968. Chafrter 5. 

Many interrelated factors affect the solubility of substances in water. This makes it challenging to predict which ionic substances will dissolve in water. By performing experiments, chemists have developed guidelines to help them make predictions about solubility. In Investigation 9-A, you will perform your own experiments to develop quidelines about the solubility of ionic compounds in water. 

How can you develop guidelines to help you predict the solubility of ionic compounds in water ... 

The solubility of ionic substances in water varies greatly. For example, sodium chloride is quite soluble in water, whereas silver chloride (contains Ag+ and Cl- ions) is only very slightly soluble. The differences in the solubilities of ionic compounds in water typically depend on the relative affinities of the ions for each other (these forces hold the solid together) and the affinities of the ions for water molecules [which cause the solid to disperse (dissolve) in water]. Solubility is a complex issue that we will explore in much more detail in Chapter 17. However, the most important thing to remember at this point is that when an ionic solid does dissolve in water, the ions are dispersed and are assumed to move around independently. 

It is more difficult to predict the solubility of polar molecular substances than to predict the solubility of ionic compounds and nonpolar molecular substances. Many polar molecular substances are soluble in both water and hexane. For example, ethanol is miscible with both water and hexane. The following generalization is helpfiil ... 

Determining fCjp from Solubility The solubilities of ionic compounds are determined experimentally, and several chemical handbooks tabulate them. Most solubility values are given in units of grams of solute dissolved in 100 grams of H2O. Because the mass of compound in solution is small, a negligible error is introduced if we assume that TOO g of water is equal to 100 mL of solution. We then convert the solubility from grams of solute per 100 mL of solution to molar solubility, the amount (mol) of solute dissolved per liter of solution (that is, the molarity of the solute). Next, we use the equation for the dissolution of the solute to find the molarity of each ion and substitute into the ion-product expression to find the value of K p. 

The solubility of ionic substances in water varies greatly. For example, sodium chloride is quite soluble in water, whereas silver chloride (contains Ag" and Cl ions) is only very slightly soluble. The differences in the solubilities of ionic compounds in water... 

Ionic compounds have very peculiar solubility trends. Some are highly soluble, whereas some others have very little solubility. The solubility of ionic compounds can be explained in terms of the interactions between the ions and the water molecules. Let s take sodium chloride as an example. Sodium chloride has a solubility of 360 g per liter or 36 g per 100 ml at room temperature. 

Assume that you have mixed two solutions, and a solid product (a precipitate) forms. How can you find out what the solid is What is its formula There are several possible approaches you can take to answering these questions. For example, we saw in Chapter 7 that we can usually predict the identity of a precipitate formed when two solutions are mixed in a reaction of this type if we know some facts about the solubilities of ionic compounds. 

Solubility is affected by temperature and by the presence of other solutes. The presence of an acid, for example, can have a major influence on the solubility of a substance. In Section 17.4 we considered the dissolving of ionic compounds in pure water. In this section we examine three factors that affect the solubility of ionic compounds (1) presence of common ions, (2) solution pH, and (3) presence of complexing agents. We will also examine the phenomenon of amphoterisniy which is related to the effects of both pH and complexing agents. 

SECTION 17.5 Several experimental factors, including temperature, affect the solubilities of ionic compounds in water. The solubility of a slightly soluble ionic compound is decreased by the presence of a second solute that furnishes a common ion (the common-ion effect). The solubility of compounds containing basic anions increases as the solution is made more acidic (as pH decreases). Salts with anions of negligible basicity (the anions of strong acids) are unaffected by pH changes. 

In order to write these equations, you have to know something about the solubility of ionic compounds. Don t fret. Here you go If a compound is soluble, it will remain in its free ion form, but if it s insoluble, it will precipitate (form a solid). Table 8-2 gives the solubilities of selected ionic compounds. 

The solubility product does not tell the entire story of the solubility of ionic compounds. The concentration of undissociated species and complex ions may be significant. In a solution of calcium sulfate, for example, about two thirds of the dissolved material dissociates to Ca and SO " and one third is undissociated CaS04(a ) (a tightly bound ion pair). In the PW2 case, species such as Pbl, Pbl2(a ), and Pbl also contribute to the total solubility. The species Pbl and PblJ are called complex ions because they are composed of simpler ions. 

The solubility of ionic compounds must be calculated from the sum of all dissolved forms. Calculation from alone yields a minimum solubility correct only if the simple ions alone are the total dissolved form, a rare case. 

The general mles for predicting the solubility of ionic compounds in water were introduced in Section 4.2. Although useful, these solubility rules do not enable us to make quantitative predictions about how much of a given ionic compound will dissolve in water. To develop a quantitative approach, we start with what we already know about chemical equilibrium. 

In this section, we explore an equilibrium system that involves the solubility of ionic compounds. Recall from Chapter 13 that most solutes, even those called soluble, have a limited solubility in a particular solvent. In a saturated solution at a particular temperature, equilibrium exists between dissolved and undissolved solute. Slightly soluble ionic compounds, which we ve been calling insoluble, reach equilibrium with very little solute dissolved. In this introductory treatment, we will assume that, as with a soluble ionic compound, the small amount of a slightly soluble ionic compound that does dissolve dissociates completely into ions. 

Soluble ionic compounds form solutions that contain many ions and therefore are strong electrolytes. To predict the solubility of ionic compounds, chemists have developed solubility rules. Table 4.1 lists eight solubiUty rules for ionic compounds. These rules apply to most of the common ionic compounds that we will discuss in Ihis course. Example 4.1 illustrates how to use the rules. 

Let ns snmmarize the main points in this section. Compounds that dissolve in water are soluble those that dissolve little, or not at all, are insoluble. Soluble substances are either electrolytes or nonelectrolytes. Nonelectrolytes form noncon-dncting aqneons solutions because they dissolve completely as molecules. Electrolytes form electrically conducting solutions in water because they dissolve to give ions in solntion. Electrolytes can be strong or weak. Almost all soluble ionic substances are strong electrolytes. Soluble molecular substances usually are nonelectrolytes or weak electrolytes the latter solution consists primarily of molecules, but has a small percentage of ions. Ammonia, NH3, is an example of a molecular substance that is a weak electrolyte. A few molecular substances (such as HCl) dissolve almost entirely as ions in the solution and are therefore strong electrolytes. The solubility rules can be used to predict the solubility of ionic compounds in water. 

Co to http //now.brookscole.com/ cracoliceSe and click Coached Problems for a tutorial on the Solubility of Ionic Compounds. 

Questions 57 and 58 In Chapter 9 we discussed how to identify major and minor species and how to write net ionic equations. These skills are based on the solubility of ionic compounds, the strengths of acids, and the stability of certain ion combinations. 

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