Chapter 16 167

Chapter 16

Acid-Base Equilibria and Solubility Equilibria

This chapter covers titration curves, buffers, and Ksp, which are all tested on the AP exam.

Students are introduced to the Henderson-Hasselbalch equation in solving buffer problems. The

equation is given to students in the AP equation page, so they are expected to be able to know

how and when to use it. However, all equilibrium problems can be solved by using combinations

of simultaneous equilibrium expressions without invoking the H-H equation.

Reading student answers often indicates that they have a very poor understanding of material in

this chapter, because they have not truly mastered the material in the previous chapter. For

example, they do not appreciate the interrelationship of conjugate acids and bases and their

respective equilibrium constants. Animations in this chapter include buffer solutions and acid

base titrations.

Major Concepts to Know

• If a base is added to a buffer system, the acid loses concentration and creates conjugate base

ions. Students should recognize every buffer system has a capacity and can be used up.

• Students can still solve buffer problems using ICE set-up, as both methods result in the same

answer.

• Understanding what pKa is will help the student later on in finding pKa from a graph.

• Students should be sure to thoroughly understand the methods used in the worked examples

in this chapter, especially examples 16.1 and 16.2. It is especially important for students to

learn and practice how to choose initial conditions so that “change” is relatively small and

approximations will work. So for instance, if

0.010 mol solid NaOH were to be added to 100 mL of a buffer solution, the initial condition

is with [OH ] = 0.10, and reaction proceeds to change from there.

• Students need to know how to prepare a buffer system of a certain pH.

• From the Henderson-Hasselbalch equation, students should know the pH is at pKa when the

conjugate base and its acid are the same molarity, so students can then estimate the pH just

by knowing Ka. A system buffers best at just ±1 pH unit from pKa (or the corresponding pOH

and pKb) because this represents a factor of 10 in concentration on each side of the equal

molarity point, allowing a “swing” of a factor of 100 in concentration.

• The concept of titrations was introduced in Chapter 4 but should be reviewed here. Students

should be able to look at any titration curve, understand the axis used, and recognize typical

acid-base reactions. Students need to recognize whether graphs are between strong acids,

weak acids, strong bases, or weak bases.

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pH Profile of a Strong-Acid Strong-Base Titration

In the above graph, students should recognize the pH starts at 1, indicating a strong acid; the

equivalence is at 7, indicating a strong acid and a strong base have reacted; and the last point

in the upper flat is 13, which also indicates a strong base.

The Relationship Between Buffer Capacity and pH Change

Two common indicators for a strong acid-strong base reaction are shown. They need to

change color at the equivalence point.

pH Profile of a Weak Acid Strong Base Titration

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In the above graph, the pH starts at 3, indicating a weak acid; the equivalence is above 7,

indicating the base is strong as also shown by the upper flat.

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pH Profile of a Strong Acid Weak Base Titration

This graph indicates a base is being titrated, and the starting pH is at 11, indicating a weak

base. The pH at equivalence is below 7, and the ending pH is 1, indicating a strong acid was

used in the titration.

• Keys to recognition are the starting pH, equivalence point position, and ending pH. Students

should be able to draw them as well. Strong acids with strong bases will have an equivalence

of 7 or neutral. Strong acids with weak bases will be in the acid range at equivalence, and

strong bases with weak acids will be basic.

• A misconception many students have is that equivalence means neutral in all cases since the

moles of H+ = moles of OH . This is a result of not understanding that the driving factor is

the hydrolysis of the salt produced by the reaction. Students need to be able to do

calculations at any point during a titration. Students also need to be able to choose

appropriate indicators given the Ka of the indicator.

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Curve for the Titration of a Weak Polyprotic Acid

• Students should recognize the different sections of a titration curve for a diprotic or triprotic

acid.

• The chapter then considers solubility of sparingly soluble substances and the resulting

equilibrium constant, Ksp. These are the same as all K problems, but since the reactant is a

solid, there is no denominator. Students need to understand when they learn a compound like

AgCl is insoluble, some ions are still in solution and an equilibrium is established.

• Confusion with vocabulary is common between molar solubility and solubility product. Use

of the calculated equilibrium quotient is best for solving prediction problems. To determine if

a precipitate will form, students need to be able to

solve for Q and then compare Q to Ksp values. Ksp values are saturation values,

so if Q > Ksp, then a precipitate will form; if Q = Ksp, the solution is saturated; and if

Q < Ksp, the solution is unsaturated and no precipitate will form. Some more complex

problems involve figuring out which of two precipitates will form first and what happens to a

concentration when a common ion is added or it has been added to an acidic or basic

solution.

• Generally, AP exam problems are relatively straightforward and do not involve multiple ions

in solution.

Chapter 16 172

• Kf problems with complex ions have not been on the AP exam. Recognizing equations for

complex ion formation has been tested as descriptive chemistry in the multiple-choice section

or in the equations question on free response.

Vocabulary to Know

• Buffer solution

• Common ion effect

• Complex ion

• End point

• Molar solubility

• Qualitative analysis

• Solubility

• Solubility product

Math Skills to Know

• pKa = log Ka

• pH = pKa + log[Conjugate base/acid]

• pOH = pKb + log[Conjugate acid/base]

• Common ion problems and determining new pH

• Buffer problems and determining new pH

• Determining pH at any point in a titration

• Solving Ksp problems

Suggested Problems (* also electronic)

• Common ion effect: 1, 2*, 3, 4, 5*, 6*

• Buffer solutions: 7, 8, 9, 10*, 11, 12, 14, 15, 17*, 18*, 20, 21, 22

• Acid-base titrations: 23, 24, 25*, 26, 27, 28*, 29*, 30, 32, 33, 35, 36

• Acid-base indicators: 37, 38, 39, 42*

• Solubility equilibria: 44, 45*, 46*, 47, 49, 50*, 51, 52, 53, 54, 55*, 56, 57*, 58

• Fractional precipitation: 59*

• Common ion effect and solubility: 61, 62*, 63, 66*

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• pH and solubility: 67*68, 70*

• Qualitative analysis: 83*, 85, 86

• Additional questions: 95, 100, 134

Suggested Demonstrations or Labs

Demos

Complexing ions demo:

o Adding NH3 to CuSO4 first forms the white/pale blue precipitate Cu(OH)2, and then, with

the addition of more concentrated NH3, the precipitate dissolves and the bright blue

complex ion, Cu(NH3)42+

, is formed.

o Adding ammonia to a solution of silver nitrate to which NaCl has been recently added

will dissolve the AgCl precipitate, forming Ag(NH3)2+ ions.

• Melanie M. Cooper, “Project 8: Buffers,” Cooperative Chemistry Lab Manual (New York,

NY: McGraw-Hill, 2006).

• Jeffrey A. Paradis, “The Properties of Buffers: Resisting Change in a Turbulent World,”

Hands On Chemistry Laboratory Manual (New York, NY: McGraw-Hill, 2006).

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Questions

(Note: *To understand this chapter, you must understand Chapter 15 on acids and bases and pH

calculations.)

Acid Base Equilibria

1. What is a common ion?

a. What does a common ion do to the solubility of a weak acid or base?

b. Explain the common ion effect.

c. Using Le Châtelier, when you add a common ion what will you shift to reduce

the stress?

2. Write the Henderson-Hasselbalch equation.

a. Explain how to get each of the values in the equation.

b. What kind of problems can the equation solve?

c. How would you rewrite the equation for a base? (not in the book but on

the AP exam)

d. Calculate the pH of a solution containing 0.20 M CH3COOH and

0.30 M NaCH3COO.

e. Calculate the pH of a solution containing 0.50 M NH3 and 0.20 NH4Cl.

3. What is a buffer?

a. How do buffers work?

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b. Explain what happens to a buffer system as an acid or a base is added.

c. What is the capacity of a buffer?

d. Calculate the pH of a buffering system containing 0.20 M CH3COOH and

0.30 M NaCH3COO when 50 mL of 0.10 M HCl is added to 100 mL of the

buffer.

e. Explain how you would make a buffer with a specific pH.

4. What does a titration curve show?

a. Draw general titration curves for reactions with (1) strong acid (SA) and

strong base (SB), (2) SA and weak base (WB), and (3) weak acid (WA)

and SB.

b. How do you find the equivalence point or end point on each?

c. What is the expected equivalence point pH of each type of titration?

d. On which curves can a Ka or Kb be determined?

• How do you find a Ka or Kb on a titration curve?

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e. If you titrate a monoprotic acid, diprotic acid, and triprotic acid, how will the

graphs of the curves be different? Include a rough sketch of each.

5. What are indicators?

a. How are they used in a titration?

b. How do you know which indicator to pick? (Explain the different types of

titrations, SA versus SB, etc.)

Solubility Equilibria

6. What are the general conditions of Ksp?

7. What is the solubility product?

a. What types of reactions use Ksp?

b. What salts don’t have Ksp values?

c. When solving a problem, how can you tell if the reaction is unsaturated,

saturated, or supersaturated?

8. What is the difference between molar solubility and solubility?

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a. How do you change from one to the other?

*Be able to do examples 16.8 and 16.9.

9. If there are two ions (s and s) present, what is the general formula for Ksp?

10. If there are three ions (2s and s) present, what is the general formula for Ksp?

11. If there are four ions (3s and s) present, what is the general formula for Ksp?

12. How can someone be able to predict if a precipitate is formed?

13. What is fractional crystallization?

a. What is fractional crystallization used for?

14. How does adding a common ion influence solubility?

a. Explain why from Le Châtelier.

15. Using Le Châtelier, explain why adding an acid to a basic solution increases the solubility of

a base.

16. Using Le Châtelier, explain why adding a base to an acidic solution increases the solubility

of an acid.

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17. What is a complex ion?

18. What metals tend to form complex ions? Why?

19. What are amphoteric hydroxides?

a. List three amphoteric hydroxides.

20. Define qualitative analysis.

a. Explain the general process of qualitative analysis.

b. Draw the flowchart from figure 16.14 on p. 756 and be able to explain what is

separated in each step.

c. How would you test for Na+ and K

+?