Chapter 17. Additional Aspects of Aqueous Equilibria
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 17.2 Buffer Action 17.2 Buffered Solutions
Figure 17.3 Calculating the pH of a Buffer After 17.2 Buffered Solutions
Addition of Acid or Base
Figure 17.7 Titration of a Strong Acid with a Strong 17.3 Acid-Base Titrations
Activities: Section:
Calculating pH Using Henderson-Hasselbalch 17.2 Buffered Solutions
Equation
Buffer pH 17.2 Buffered Solutions
Animations: Section:
Common-Ion Effect 17.1 The Common-Ion Effect
Acid-Base Titration 17.3 Acid-Base Titrations
Dissolution of Mg(OH)2 by Acid 17.5 Factors That Affect Solubility
Chapter 17
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3-D Models: Section:
Acetic Acid 17.1 The Common-Ion Effect
Hydroxide Ion 17.3 Acid-Base Titrations
Other Resources
Further Readings: Section:
A Good Idea Leads to a Better Buffer 17.2 Buffered Solutions
Phosphate Buffers and Telephone Poles—A Useful 17.2 Buffered Solutions
Analogy with Limitations
Assessing Students’ Conceptual Understanding of 17.4 Solubility Equilibria
Solubility Equilibrium
What Should We Teach Beginners about Solubility 17.4 Solubility Equilibria
and Solubility Products?
Chemical Aspects of Dentistry 17.5 Factors That Affect Solubility
Dentifrice Fluoride 17.5 Factors That Affect Solubility
Live Demonstrations: Section:
The Common Ion Effect: Second Demonstration 17.1 The Common-Ion Effect
Effect of Acetate Ion on the Acidity of Acetic Acid: 17.1 The Common-Ion Effect
Additional Aspects of Aqueous Equilibria
247
Buffering Action and Capacity 17.2 Buffered Solutions
Buffering Action of Alka-Seltzer 17.2 Buffered Solutions
Equilibrium: The Dissociation of Acetic Acid 17.2 Buffered Solutions
Determination of Neutralizing Capacity of Antacids 17.2 Buffered Solutions
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Chapter 17. Additional Aspects of Aqueous Equilibria
Common Student Misconceptions
• Students often believe that the pH at the equivalence point for any titration is 7.00; in other words,
students often think that neutralization always results in the formation of a neutral solution.
• Students often think that titration is a new type of a reaction, rather than an experimental technique.
Teaching Tips
• Students should review Le Châtelier’s principle, pH of salts, and solubility rules prior to starting this
chapter.
Lecture Outline
17.1 The Common Ion Effect
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• The dissociation of a weak electrolyte is decreased by the addition of a strong electrolyte that has an
ion in common with the weak electrolyte.
• For example, consider the ionization of a weak acid, acetic acid.
HC2H3O2(aq) H+(aq) + C2H3O2–(aq)
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“The Common–Ion Effect: Second Demonstration” from Live Demonstrations
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“Common–Ion Effect” Animation from Instructor’s Resource CD/DVD
Additional Aspects of Aqueous Equilibria
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17. 2 Buffered Solutions
• A buffered solution, or buffer, is a solution that resists a drastic change in pH upon addition of small
amounts of strong acid or strong base.
Composition and Action of Buffered Solutions
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• A buffer consists of a mixture of a weak acid (HX) and its conjugate base (X– ).
HX(aq) H+(aq) + X– (aq)
Calculating the pH of a Buffer
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• The pH of the buffer is related to Ka and to the relative concentrations of the acid and base.
• We can derive an equation that shows the relationship between conjugate acid-base concentrations,
pH and Ka.
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“A Good Idea Leads to a Better Buffer” from Further Readings
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“Phosphate Buffers and Telephone Poles—A Useful Analogy with Limitations” from Further Readings
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“Buffering Action and Capacity” from Live Demonstrations
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“Buffering Action of Alka–Seltzer” from Live Demonstrations
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• An alternate form of this equation is:
Buffer Capacity and pH Range
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• Buffer capacity is the amount of acid or base that can be neutralized by the buffer before there is a
significant change in pH.
• Buffer capacity depends on the concentrations of the components of the buffer.
• The greater the concentrations of the conjugate acid-base pair, the greater the buffer capacity.
• The pH range of a buffer is the pH range over which it is an effective buffer.
• The pH range of a buffer is generally within one pH unit of the pKa of the buffering agent.
Addition of Strong Acids or Bases to Buffers
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• Keep in mind that reactions between strong acids and weak bases proceed essentially to completion.
• The same is true for reactions between strong bases and weak acids.
• If we do not exceed the buffering capacity of the buffer, then the added strong acid or base is
completely consumed by reaction with the buffer.
17.3 Acid-Base Titrations
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• In an acid-base titration:
• a solution of base of known concentration is added to an acid (or an acid of known concentration
is added to a base).
• acid-base indicators, or pH meters, are used to signal the equivalence point.
• The equivalence point is the point at which stoichiometrically equivalent quantities of acid
and base have been added.
• The plot of pH versus volume during a titration is called a pH titration curve.
Additional Aspects of Aqueous Equilibria
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Strong Acid–Strong Base Titrations
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• Consider adding a strong base (e.g., NaOH) to a solution of a strong acid (e.g., HCl).
• We can divide the titration curve into four regions.
• 1. Initial pH (before any base is added)
• The pH is given by the strong acid solution.
• Therefore, pH < 7.
• The shape of a strong base-strong acid titration curve is very similar to a strong acid-strong base
titration curve.
• Initially, the strong base is in excess, so the pH > 7.
• As acid is added, the pH decreases but is still greater than 7.
• At the equivalence point, the pH is given by the salt solution (i.e., pH = 7).
• After the equivalence point, the pH is given by the strong acid in excess, so pH is less than 7.
Weak Acid-Strong Base Titration
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• Consider the titration of acetic acid, HC2H3O2, with NaOH.
• Again, we divide the titration into four general regions:
• 1. Before any base is added.
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“Acid–Base Indicators: A New Look at an Old Topic” from Further Readings
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“Teas as Natural Indicators” from Live Demonstrations
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“Acid–Base Titration” Activity from Instructor’s Resource CD/DVD
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“Edible Acid–Base Indicators” from Further Readings
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“Predicting Acid–Base Titration Curves without Calculations” from Further Readings
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• 4. After the equivalence point.
• the pH is given by the concentration of the excess strong base.
• The pH curve for a weak acid-strong base titration differs significantly from that of a strong acid-
strong base titration.
• For a strong acid-strong base titration:
• the pH begins at less than 7 and gradually increases as base is added.
• Near the equivalence point, the pH increases dramatically.
• For a weak acid-strong base titration:
basic salt.
Titrations of Polyprotic Acids
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• In polyprotic acids, the ionizable protons dissociate in a series of steps.
• Therefore, in a titration there are n equivalence points corresponding to each ionizable proton.
• In the titration of H3PO4 with NaOH there are three equivalence points:
• one for the formation of H2PO4–,
• one for the formation of HPO42–, and
• one for the formation of PO43–.
Titrating with an Acid-Base Indicator
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• How can we analyze the titration (i.e., how will we know when we are at the equivalence point)?
• We often use a pH indicator.
Additional Aspects of Aqueous Equilibria
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• The pH changes rapidly only over the pH range from about pH 11 to 3.
• To detect the equivalence point, we use an indicator that changes color somewhere close to
the pH at the equivalence point.
• Usually, we use phenolphthalein, which changes color between pH 8.3 and 10.0.
FORWARD REFERENCES
• Redox titrations will be briefly mentioned in Chapter 20 (section 20.2).
• Oxyacids of phosphorus will be discussed in more detail in Chapter 22 (section 22.8).
17.4 Solubility Equilibria
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The Solubility-Product Constant, Ksp
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• Consider a saturated solution of BaSO4 in contact with solid BaSO4.
• We can write an equilibrium expression for the dissolving of the slightly soluble solid.
BaSO4(s) Ba2+(aq) + SO42–(aq)
Solubility and Ksp
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• Solubility is the amount of substance that dissolves to form a saturated solution.
• This is often expressed as grams of solute that will dissolve per liter of solution.
• Molar solubility is the number of moles of solute that dissolve to form a liter of saturated solution.
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“Solubility of Some Silver Compounds” from Live Demonstrations
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“Red and White Precipitates in Sodium Silicate” from Live Demonstrations
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“The Murky Pool” from Further Readings
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• We can use the solubility to find Ksp and vice versa.
• To convert solubility to Ksp:
• Convert solubility into molar solubility (via molar mass).
FORWARD REFERENCES
• The role of acid-base and solubility reactions in tying the ocean to the atmosphere and the
global climate will be discussed in Chapter 18 (sections 18.3 and 18.4).
• Water softening to remove hardness ions will be discussed in Chapter 18 (section 18.4).
• Calculations involving Ksp versus Grxno will be performed in Chapter 19 (section 19.7).
17.5 Factors That Affect Solubility
• Three factors that have a significant impact on solubility are:
• The presence of a common ion,
Common-Ion Effect
• The solubility of a slightly soluble salt is decreased when a common ion is added.
• This is an application of Le Châtelier’s principle.
• Consider the solubility of CaF2:
CaF2(s) Ca2+(aq) + 2F–(aq)
produces a common ion.
Solubility and pH
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• Again, we apply Le Châtelier’s principle:
Mg(OH)2(s) Mg2+(aq) + 2OH–(aq)
• If OH– is removed, then the equilibrium shifts toward the right and Mg(OH)2 dissolves.
• OH– can be removed by adding a strong acid:
OH–(aq) + H+(aq) H2O(aq)
Additional Aspects of Aqueous Equilibria
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PbF2(s) Pb2+(aq) + 2F–(aq)
• If the F– is removed, then the equilibrium shifts towards the right and PbF2 dissolves.
• F– can be removed by adding a strong acid:
Formation of Complex Ions
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• Recall that metal ions may act as Lewis acids in aqueous solution (water may act as the Lewis base).
• Such an interaction may have a significant impact on metal salt solubility.
• For example, AgCl has a very low solubility.
• Ksp for AgCl = 1.8 10–10
• However, the solubility is greatly increased if ammonia is added.
• Why?
• Consider the formation of Ag(NH3)2+:
Ag+(aq) + 2NH3(aq) Ag(NH3)2+(aq)
• The Ag(NH3)2+ is called a complex ion.
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“Chemical Aspects of Dentistry” from Further Readings
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“Dentifrice Fluoride” from Further Readings
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“Fluorine Compounds and Dental Health: Applications of General Chemistry Topics” from Further
Readings
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Amphoterism
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• Substances that are capable of acting either as an acid or a base are amphoteric.
• The term is similar to one discussed earlier: amphiprotic, which relates more generally to any
species that can either gain or lose a proton.
• Amphoteric metal oxides and hydroxides will dissolve in either a strong acid or a strong base.
• Examples are hydroxides and oxides of Al3+, Cr3+, Zn2+, and Sn2+.
• The hydroxides generally form complex ions with several hydroxide ligands attached to the
FORWARD REFERENCES
• Chelating agents will be further mentioned in Chapters 23 (section 23.3, respectively).
• Amphoteric oxides of transition metals will be mentioned in Chapter 22 (section 22.5).
17.6 Precipitation and Separation of Ions
• Consider the following:
BaSO4(s) Ba2+(aq) + SO42–(aq)
• At any instant in time, Q = [Ba2+][SO42–].
• If Q > Ksp, precipitation occurs until Q = Ksp.
• If Q = Ksp equilibrium exists (saturated solution).
• If Q < Ksp, solid dissolves until Q = Ksp.
Selective Precipitation of Ions
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• Ions can be separated from each other based on the solubilities of their salts.
• Example: If HCl is added to a solution containing Ag+ and Cu2+, the silver precipitates (Ksp for
AgCl is 1.8 10–10) while the Cu2+ remains in solution.
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“Precipitation Reactions” Movie from Instructor’s Resource CD/DVD
Additional Aspects of Aqueous Equilibria
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• When more H2S is added to the solution, a second precipitate of white ZnS forms.
17.7 Qualitative Analysis for Metallic Elements
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• Quantitative analysis is designed to determine how much metal ion is present.
• Qualitative analysis is designed to detect the presence of metal ions.
• Typical qualitative analysis of a metal ion mixture involves:
• 1. separation of ions into five major groups on the basis of their differential solubilities.
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Further Readings:
1. Charles L. Bering, “A Good Idea Leads to a Better Buffer,” J. Chem. Educ., Vol. 64, 1987, 803–805.
2. Edwin S. Gould, “Phosphate Buffers and Telephone Poles—A Useful Analogy with Limitations,” J.
Chem. Educ., Vol. 76, 1999, 1511.
7. Robert C. Mebane and Thomas R. Rybolt, “Edible Acid–Base Indicators,” J. Chem. Educ., Vol. 62,
1985, 285.
8. Dennis Barnum, “Predicting Acid-Base Titration Curves without Calculations,” J. Chem. Educ., Vol.
76, 1999, 938–942.
9. Robert Perkins, “The Useless Tea Kettle,” J. Chem. Educ., Vol. 61, 1984, 383.
10. Roy W. Clark and Judith M. Bonicamp, “The Ksp–Solubility Conundrum,” J. Chem. Educ., Vol. 75,
1998, 1182–1185.
11. Robert Perkins, “The Murky Pool,” J. Chem. Educ., Vol. 61, 1984, 383–384.
Additional Aspects of Aqueous Equilibria
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18. Kim Chwee, Daniel Tan, Ngoh Khang Goh, Lian Sai Chai, and David F. Teagust, “Major Sources of
Difficulty in Students’ Understanding of Basic Inorganic Qualitative Analysis,” J. Chem. Educ., Vol. 81,
2004, 725–732.
Live Demonstrations:
1. Lee. R. Summerlin, and James. L. Ealy, Jr., “The Common Ion Effect: Second Demonstration,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 1 (Washington: American Chemical
Society, 1988), pp. 93–94. The reaction of calcium carbonate and acetic acid is used to demonstrate the
common ion effect.
5. Bassam Z. Shakhashiri, “Buffering Action of Alka–Seltzer,” Chemical Demonstrations: A Handbook
for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989), pp. 186–187.
6. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Equilibrium: The Dissociation of Acetic
Acid,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American
Chemical Society, 1988), pp.160–161. Changes in indicator color upon addition of base or acetate to
acetic acid are explored.
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10. Lee R. Summerlin, Christie L Borgford, and Julie B. Ealy, “Silver Ion Solubilities: Red and White
Precipitates,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American
Chemical Society, 1988), pp. 124–125. This is an effective introduction to equilibrium; the relative
solubilities of silver chromate and silver chloride are investigated.
13. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “ Colorful Complex Ions in Ammonia,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), pp. 75–76. Ammine complexes of copper and cobalt are prepared in this demonstration.
14. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “ Green and Blue Copper Complexes,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), pp.71–72. Three copper complexes are prepared in this demonstration.