Ionic compounds solubility rules

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. 

The solubihty rules are an empirical set of guidelines that help predict the solubilities of ionic compounds these rules are especially useful when determining whether or not a precipitate wiU form. 

STRATEGY Decide which ions are present in the mixed solutions and consider all possible combinations. Use the solubility rules in Table 1.1 to decide which combination corresponds to an insoluble compound and write the net ionic equation to match. 

The general criterion for solubility is the rule that like dissolves like . In other words polar solvents dissolve polar and ionic solutes, non-polar solvents dissolve non-polar solutes. In the case of water, this means that ionic compounds such as sodium chloride and polar compounds such as sucrose are soluble, but non-polar compounds such as paraffin wax are not. 

In each case, each available cation is paired with the available anions, one at a time, to determine if a compound is produced that is insoluble, based on the solubility rules of Chapter 5. Then a net ionic equation is written to summarize this information. 

These problems involve mixing two solutions. Each solution is a water solution of an ionic compound. From the mixture of the two solutions, at least one insoluble precipitate will form. The other ions present are probably soluble and are called spectator ions they are not included in the net ionic equation. You must know your solubility rules to do these problems. 

Showing the strong electrolytes in the form of ions yields the ionic equation (sometimes called the total or overall ionic equation). The strong electrolytes are any strong acid, strong base, or water soluble (according to the solubility rules) ionic compound. In this example, the ionic equation is ... 

TABLE 2.1 Solubility Rules for Ionic Compounds in Water... 

Water. It should come as no surprise that ordinary water can be an excellent solvent for many samples. Due to its extremely polar nature, water will dissolve most substances of likewise polar or ionic nature. Obviously, then, when samples are composed solely of ionic salts or polar substances, water would be an excellent choice. An example might be the analysis of a commercial iodized table salt for sodium iodide content. A list of solubility rules for ionic compounds in water can be found in Table 2.1. 

The method makes use of only those species, dissolved or otherwise, that actually take part in the reaction. So-called spectator ions, or ions that are present but play no role in the chemistry, are not included in the balancing procedure. Solubility rules are involved here, since spectator ions result only when an ionic compound dissolves and ionizes. Also, the scheme is slightly different for acid and base conditions. Our purpose, however, is to discuss the basic procedure thus spectator ions will be absent from all examples from the start, and acidic conditions will be the only conditions considered. The stepwise procedure we will follow is below ... 

Use the Ionic Compounds activity (eChapter 4.4) to determine the formula of each of the insoluble iron(III) salts. Then, using your knowledge of the solubility rules, write a molecular, ionic, and net ionic equation for an aqueous reaction that would produce each salt. 

A mixed convention will be used in this chapter as an aid to indicate whether a given compound may be written in the ionized form. We will use Na+Cl-, Ba2+(NO )2, and similar notation, to indicate that compounds are ionic. Naturally, the compounds could be expressed by the molecular formulas, NaCl and Ba(N03)2, which works well for those who are intimately familiar with the solubility rules. We are not presenting the solubility rules since we prefer to concentrate on balancing the equations without having the distraction of investigating the solubility rules at the same time. 

One of the great rules of thumb in chemistry is like dissolves like. That means that polar solutes are more soluble in polar solvents, while nonpolar solutes are more soluble in nonpolar solvents. We have discussed molecular polarity in a previous chapter, but you can consider most organic compounds to be nonpolar. The most polar species are, obviously, ionic compounds, followed by species that can form hydrogen bonds, such as water and ethanol (CH3CH2OH). Therefore, you would expect ionic compounds to be soluble in water, but not very soluble in an organic solvent such as ether or hexane. On the other hand, you would expect an organic compound, like the vast majority of the pure form of injectable medications, to be relatively insoluble in a water-based medium such as blood. 

Next, the ionic equation is written. The ionic equations show the strong electrolytes (soluble compounds as predicted by the solubility rules) as ions. Of course, solid compounds are not separated. 

Ionic equation Shows the strong electrolytes (soluble compounds as predicted by the solubility rules) as ions. 

In aqueous solutions of ionic compounds, the ions act independently of each other. Soluble ionic compounds are written as their separate ions. We must be familiar with the solubility rules presented in Chapter 8 and recognize that the following types of compounds are strong electrolytes strong acids in solution, soluble metallic hydroxides, and salts. (Salts, which can be formed as the products of reactions of acids with bases, include all ionic compounds except strong acids and bases and metalhc oxides and hydroxides.) Compounds must be both ionic and soluble to be written in the form of their separate ions. (Section 9.1)... 

The solubility characteristics of non-ionic compounds are determined chiefly by their polarity. Non-polar or weakly polar compounds dissolve in non-polar or weakly polar solvents highly polar compounds dissolve in highly polar solvents. Like dissolves like is an extremely useful rule of thumb. Methane dissolves in carbon tetrachloride because the forces holding methane molecules to each other and carbon tetrachloride molecules to each other are replaced by very similar forces holding methane molecules to carbon tetrachloride molecules. 

Table 2 Solubility Rules for Some Common Ionic Compounds...

We can assume that the precipitate that formed in our test tube must be one (or both) of the possible products from our word equation. In order to identify which of the products formed, we would need to find out which of the products is insoluble in water and would therefore precipitate out. Our most handy reference tables for this type of problem will contain the solubility rules for ionic compounds, shown here, and/or a solubility table (see Figure 6-3a). 

Notice that the title of the rules is general solubility rules. There are exceptions to the rules listed there. For example, although ionic compounds that contain halogens tend to be soluble, when the cation is lead (II) (as in lead (II) chloride or lead (II) iodide) the ionic compound will be insoluble in water. The rules listed are really meant to be general trends, which you may want to memorize. They will give you the ability to make fast predictions about the identity of a precipitate that is found after many ionic reactions, just not the one in our first example. 

The factors that determine whether or not an ionic compound will be soluble in water are complex, making predictions difficult. As a result, a series of statements or rules have come into being that guide predictions. These rules are called the solubility rules. As with most general rules, exceptions exist, but they are correct most of the time. The following rules are organized in a hierarchal structure—that is, the first rule takes precedence over the second, the second over the third, and so forth. 

Let s apply these rules to predict the solubility of a few ionic compounds. 

As a general rule, the larger the charges on the cation and anion of an ionic compound, the less likely will the compound be soluble in water, especially if both ions carry multiple charges, (page 360)... 

TABLE 4.2 Solubility Rules for Common Ionic Compounds in Water at 25°C... 

From general rules about solubilities of ionic compounds, we can predict whether a precipitate will form in a reaction. 

The rules given in Table 4.2 (p. 113) allow us to predict the solubility of a particular ionic compound in water. When sodium chloride dissolves in water, the ions are stabilized in solution by hydration, which involves ion-dipole interaction. In general, we predict that ionic compounds should be much more soluble in polar solvents, such as water, liquid ammonia, and liquid hydrogen fluoride, than in nonpolar solvents, such as benzene and carbon tetrachloride. Since the molecules of nonpolar solvents lack a dipole moment, they cannot effectively solvate the Na and Cl ions. Solvation is the process in which an ion or a molecule is surrounded by solvent molecules arranged in a specific manner. The process is called hydration when the solvent is water.) The predominant intermolecular interaction between ions and nonpolar compounds is ion-induced dipole interaction, which is much weaker than ion-dipole interaction. Consequently, ionic compounds usually have extremely low solubility in nonpolar solvents. 

In this section, we explore the aqueous equilibria of slightly soluble ionic compounds, which up to now we ve called insoluble. In a saturated solution at a particular temperature, equilibrium exists between the undissolved and dissolved solute (Chapter 13). Slightly soluble ionic compounds have a relatively low solubility, so they reach equilibrium with relatively little solute dissolved. At this point, it would be a good idea for you to review the solubility rules listed in Table 4.1. 

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