# FOCUS KIMIA

## Diluting Concentrated Solutions

For convenience, chemicals are sometimes bought and stored as concentrated solutions, which are then diluted before use. Aqueous hydrochloric acid, for example, is sold commercially as a 12.0 M solution, yet it is most commonly used in the laboratory after dilution with water to a final concentration of either 6.0 M or 1.0 M.

Concentrated solution Solvent → Dilute solution

The main thing to remember when diluting a concentrated solution is that the number of moles of solute is constant; only the volume of the solution is changed by adding more solvent. Because the number of moles of solute can be calculated by multiplying molarity times volume, we can set up the following equation:

where Mi is the initial molarity, Vi is the initial volume, Mf is the final molarity, and Vf is the final volume after dilution. Rearranging this equation into a more useful form shows that the molar concentration after dilution (Mf) can be found by multiplying the initial concentration (Mi) by the ratio of initial and final volumes (Vi > Vf):

## Precipitation Reactions in Aqueous Solution

One common type of reaction that occurs in aqueous solution is the precipitation reaction, which results in the formation of an insoluble product, or precipitate. A precipitate is an insoluble solid that separates from the solution. Precipitation reactions usually involve ionic compounds. For example, when an aqueous solution of lead(II) nitrate [Pb(NO3)2] is added to an aqueous solution of potassium iodide (KI), a yellow precipitate of lead iodide (PbI2) is  formed:

Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

Potassium nitrate remains in solution. Figure 1 shows this reaction in progress. The preceding reaction is an example of a metathesis reaction (also called a double displacement reaction), a reaction that involves the exchange of parts between two compounds. (In this case, the compounds exchange the NO3- and I- ions.) As we will see, the precipitation reactions discussed in this chapter are examples of metathesis reactions.

Figure 1 Formation of yellow PbI2 precipitate as a solution of Pb(NO3)2 is added to a solution of KI.

Solubility
How can we predict whether a precipitate will form when a compound is added to a solution or when two solutions are mixed? It depends on the solubility of the solute, which is defned as the maximum amount of solute that will dissolve in a given quantity of solvent at a specif  c temperature. Chemists refer to substances as soluble, slightly soluble, or insoluble in a qualitative sense. A substance is said to be soluble if a fair amount of it visibly dissolves when added to water. If not, the substance is described as slightly soluble or insoluble. All ionic compounds are strong electrolytes, but they are not equally soluble. Table 1 classif es a number of common ionic compounds as soluble or insoluble. Keep in mind, however, that even insoluble compounds dissolve to a certain extent. Figure 2 shows several precipitates.

Table 1 Solubility Rules for Common Ionic Compounds in Water at 25˚C
Molecular Equations, Ionic Equations, and Net Ionic Equations
The equation describing the precipitation of lead iodide on page 100 is called a molecular equation because the formulas of the compounds are written as though all species existed as molecules or whole units. A molecular equation is useful because it identifes the reagents (that is, lead nitrate and potassium iodide). If we wanted to bring about this reaction in the laboratory, we would use the molecular equation. However, a molecular equation does not describe in detail what actually is happening in solution.

Figure 2 Appearance of several precipitates. From left to right: CdS, PbS, Ni(OH)2, Al(OH)3.

As pointed out earlier, when ionic compounds dissolve in water, they break apart into their component cations and anions. To be more realistic, the equations should show the dissociation of dissolved ionic compounds into ions. Therefore, returning to the reaction between potassium iodide and lead nitrate, we would write

## Explanation About Electrolytes versus Nonelectrolytes

All solutes that dissolve in water fit into one of two categories: electrolytes and nonelectrolytes. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte does not conduct electricity when dissolved in water. Figure 1 shows an easy and straightforward method of distinguishing between electrolytes and nonelectrolytes. A pair of platinum electrodes is immersed in a beaker of water. To light the bulb, electric current must ﬂow from one electrode to the other, thus completing the circuit. Pure water is a very poor conductor of electricity. However, if we add a small amount of sodium chloride (NaCl), the bulb will glow as soon as the salt dissolves in the water. Solid NaCl, an ionic compound, breaks up into Na+ and Cl- ions when it dissolves in water. The Na+ ions are attracted to the negative electrode and the Cl- ions to the positive electrode. This movement sets up an electrical current that is equivalent to the ﬂow of electrons along a metal wire. Because the NaCl solution conducts electricity, we say that NaCl is an electrolyte. Pure water contains very few ions, so it cannot conduct electricity. Comparing the lightbulb’s brightness for the same molar amounts of dissolved substances helps us distinguish between strong and weak electrolytes. A characteristic of strong electrolytes is that the solute is assumed to be 100 percent dissociated into ions in solution. (By dissociation we mean the breaking up of the compound into cations and anions.) Thus, we can represent sodium chloride dissolving in water as 