Book Cover Chemistry 8th Edition / Chang
Student Study Guide

Chapter 16: Acid-Base Equilibria and Solubility Equilibria


Index | 16.2 – 16.3 | 16.4 – 16.5 | 16.6 | 16.6 – 16.7 | 16.8 – 16.9 | 16.10 |

TITRATION CURVES AND INDICATORS (16.4 – 16.5)

STUDY OBJECTIVES

  1. Describe the shape of titration curves for titrations involving a strong acid and a strong base, a weak acid and a strong base, and a strong acid and a weak base.
  2. Choose the best indicator for a particular acid-base titration.

Titration Curves. Acid-base titrations were discussed in Chapter 4. With the introduction to pH in the previous chapter, it is informative to follow the pH as a function of the progress of the titration. A graph of pH versus volume of titrant added is called a titration curve. Initially, the pH is that of the unknown solution. As titrant is added, the pH becomes that of a partially neutralized solution of unknown plus titrant. The pH at the equivalence point refers to the H+ ion concentration when just enough titrant has been added to completely neutralize the unknown. If more titrant is added after the equivalence point has been reached, the pH assumes a value consistent with the pH of excess titrant. We will review three types of titration curves.

Strong acid-strong base titrations. visual aid Figure 16.3 of the text shows a titration curve for the addition of a strong base to a strong acid. The main features of this curve can be stated briefly. 1. The pH starts out quite low because this is the pH of the pure acid solution. 2. As base is slowly added, it is neutralized, and the pH is determined by the unreacted excess acid. Near the equivalence point, the pH begins to rise more rapidly. 3. At the equivalence point the pH changes sharply, increasing about 5.0 units upon the addition of only two drops of base. 4. Beyond the equivalence point the pH is determined by the amount of excess base that is added and is higher than 7.

The pH at the equivalence point of an acid-base titration is the pH of the salt solution that is formed by neutralization. NaCl, NaNO3, NaBr, KCl, and KI are examples of salts that can be formed in titrations of strong acids and strong bases. These salts yield ions that do not cause hydrolysis (Chapter 15). Therefore, the pH at the equivalence point in a titration of a strong acid with a strong base is 7.

Weak acid-strong base titrations. visual aid Figure 16.4 of the text shows a titration curve for the addition of a strong base to a weak acid. Because the acid is a weak acid, the initial pH is greater than in the titration of a strong acid. At the equivalence point, the pH is above 7.0 because of salt hydrolysis. The anion of the salt is the conjugate base of the weak acid used in the titration. In the titration of acetic acid with sodium hydroxide the salt produced is CH3COONa. The Na+ ion does not hydrolyze, but CH3COO– does:

CH3COO– + H2O CH3COOH + OH–

Hydrolysis of acetate ions makes the solution basic at the equivalence point.

Strong acid-weak base titrations. visual aid Figure 16.5 of the text shows a titration curve for the addition of a weak base to a strong acid. The first part of the curve is the same as in the strong acid versus strong base titration. However, the pH at the equivalence point is below 7.0 because the cation of the salt is the conjugate acid of the weak base used in the titration. Hydrolysis of this salt yields an acidic solution. In the titration of hydrochloric acid with ammonia, the salt produced is NH4Cl. The Cl– ion does not hydrolyze, but does:

NH4+ + H2O NH3 + H3O+

 


EXAMPLE Net Ionic Equations

Write the net ionic equations for the neutralization reactions that occur during the following titrations. Predict whether the pH at the equivalence point will be above, below, or equal to 7.0.

a. Titration of HI with NH3.

no reaction and pH
HI + NH3 NH4+ + I  
H+ + NH3 NH4+
I- + NH3 NH3I-

b. Titration of HI with NaOH.

no reaction and pH
HI + NaOH H2O + NaI  
HI + OH- H2O + I-
H+ + NaOH H2O + Na+
H+ + OH- H2O  
I- + Na+ NaI
I- + NaOH NaI + OH-

c. Titration of HF with NaOH.

no reaction and pH
HF + NaOH H2O + NaF  
HF + OH- H2O + F-
H+ + NaOH H2O + Na+
H+ + OH- H2O  
F- + Na+ NaF
F- + NaOH NaF + OH-


Indicators. Indicators are used in the laboratory to reveal the equivalence point of a titration. The abrupt change in color of the indicator signals the endpoint of a titration which usually coincides very nearly with the equivalence point.

Indicators are usually weak organic acids that have distinctly different colors in the nonionized (molecular) and ionized forms.

     

To determine the pH range in which an indicator will change color we can write the Ka expression in logarithmic form.

The color of an indicator depends on which form predominates. Typically, when [In–]/[HIn] 10 the solution will be color 2, and when [In–]/[HIn] 0.1, the solution will be color 1. Thus the color change will occur between

pH = pKa + log (10/1) = pKa + 1.0

and

pH = pKa + log (1/10) = pKa – 1.0

At the midpoint of the pH range over which the color changes, [HIn] = [In–], and pH = pKa.

Like any weak acid each HIn has a characteristic pKa, and so each indicator changes color at a characteristic pH. visual aid Table 16.1 of the text lists a number of indicators used in acid-base titration and the pH ranges over which they change color.

The choice of indicator for a particular titration depends on the expected pH at the equivalence point. For instance, in the titration of acetic acid by sodium hydroxide, which is discussed in Example 16.5 of the text, the pH at the equivalence point is 8.72. According to Table 16.1 both cresol red and phenolphthalein change color over ranges that include pH 8.72. Therefore either of these indicators would show the equivalence point of this titration.


EXAMPLE Choosing an Indicator

Choose an indicator for the titration of 50 mL of a 0.10 M HI solution with 0.10 M NH3. Consult a visual aid list of indicators.

HI(aq) + NH3(aq) NH4I(aq)

or

pH =


OBJECTIVE CHECK

Complete the following questions to check your understanding of the material. Select the check button to see if you answered correctly.

6. Approximately what range of pH should be expected at the equivalence points in the titration of weak acids with strong bases?

7. The pH at the equivalence point in a titration was found to be approximately 5. Give the category of titration in terms of the acid and base strengths.

8. 25.0 mL of 0.222 M HBr was titrated with 0.111 M NaOH.

a. Write the overall balanced equation for reaction.

b. After adding 30.0 mL of base solution, what is the pH?

9. Calculate the pH in the following titration after the addition of 12.0 mL of 0.100 M KOH to 20.0 mL of 0.200 M CH3COOH.

10. Determine the pH at the equivalence point in the titration of 0.100 M NaOH with 0.100 M HNO3.



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