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Chapter 16. Acid-Base Equilibria
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 16.3 Relative Strengths of Select Conjugate 16.2 Brønsted-Lowry Acids and Bases
Acid-Base Pairs
Figure 16.5 H+ Concentration of H+ and pH Values 16.4 The pH Scale
of Some Common Substances at 25 °C
Activities: Section:
Conjugate Acids and Bases 16.2 Brønsted-Lowry Acids and Bases
Kw 16.3 The Autoionization of Water
pH Estimation 16.4 The pH Scale
Acids and Bases 16.4 The pH Scale
Equilibrium Constant 16.6 Weak Acids
Animations: Section:
Introduction to Aqueous Acids 16.5 Strong Acids and Bases
Movies: Section:
Natural Indicators 16.4 The pH Scale
3-D Models: Section:
Oxalic Acid Introduction
Water 16.2 Brønsted-Lowry Acids and Bases
Hydronium Ion 16.2 Brønsted-Lowry Acids and Bases
Hydrogen Chloride 16.2 Brønsted-Lowry Acids and Bases
Chapter 16
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Phosphorus Pentachloride 16.11 Lewis Acids and Bases
Other Resources
Further Readings: Section:
Acids and Bases 16.1 Acids and Bases: A Brief Review
The Origin of the Term ‘Base’ 16.1 Acids and Bases: A Brief Review
Do pH in Your Head 16.4 The pH Scale
Teaching the Truth about pH 16.4 The pH Scale
The pH Concept 16.4 The pH Scale
Defining and Teaching pH 16.4 The pH Scale
The Symbol for pH 16.4 The pH Scale
The Correlation of Binary Acid Strengths with 16.10 Acid-Base Behavior and Chemical
Molecular Properties in First-Year Chemistry Structure
The Chemistry of Swimming Pool Maintenance 16.10 Acid-Base Behavior and Chemical
Structure
The Research Style of Gilbert N. Lewis: Acids and 16.11 Lewis Acids and Bases
Bases
Live Demonstrations: Section:
Food is Usually Acidic, Cleaners Are Usually Basic 16.4 The pH Scale
Colorful Acid-Base Indicators 16.4 The pH Scale
Acid-Base Equilibria
231
Chapter 16. Acid-Base Equilibria
Common Student Misconceptions
• Students often confuse a weak acid with a dilute acid.
• Students have problems with the numerical parts of this chapter. They should be strongly encouraged
Teaching Tips
• Determining pH of salts is often very conceptually and mathematically challenging. Solving and
assigning many problems is needed to overcome these difficulties.
• Students should be made aware of the fact that terms such as hydrogen ion, aqueous proton – or
Lecture Outline
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16.1 Acids and Bases: A Brief Review
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• Acids taste sour and cause certain dyes to change color.
• Bases taste bitter and feel soapy.
• Arrhenius concept of acids and bases:
16.2 Brønsted-Lowry Acids and Bases
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• We can use a broader, more general definition for acids and bases that is based on the fact that acid-
base reactions involve proton transfers.
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“Oxalic Acid” 3-D Model from Instructor’s Resource CD/DVD
Chapter 16
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The H+ Ion in Water
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• The H+(aq) ion is simply a proton (nucleus of a hydrogen atom without its valence electron).
• In water, clusters of hydrated H+(aq) ions form.
• The simplest cluster is H3O+(aq).
• We call this a hydronium ion.
• Larger clusters are also possible (such as H5O2+ and H9O4+).
• Generally, we use H+(aq) and H3O+(aq) interchangeably.
Proton-Transfer Reactions
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• We will focus our attention on H+(aq).
• According to the Arrhenius definitions, an acid increases [H+] and a base increases [OH–].
• Another definition of acids and bases was proposed by Brønsted and Lowry.
• In the Brønsted-Lowry system, a Brønsted-Lowry acid is a species that donates H+ and a Brønsted-
Lowry base is a species that accepts H+.
• Therefore, a Brønsted-Lowry base does not need to contain OH–.
Conjugate Acid-Base Pairs
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• Whatever is left of the acid after the proton is donated is called its conjugate base.
• Similarly, a conjugate acid is formed by adding a proton to the base.
• Consider HX(aq) + H2O(l) H3O+(aq) + X–(aq):
Relative Strengths of Acids and Bases
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• The stronger an acid is, the weaker its conjugate base will be.
• We can categorize acids and bases according to their behavior in water.
• 1. Strong acids completely transfer their protons to water.
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“Water” 3-D Model from Instructor’s Resource CD/DVD
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“Hydronium Ion” 3-D Model from Instructor’s Resource CD/DVD
Acid-Base Equilibria
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• No undissociated molecules remain in solution.
• Their conjugate bases have negligible tendencies to become protonated.
• An example is HCl.
• H+ is the strongest acid that can exist in equilibrium in aqueous solution.
• OH– is the strongest base that can exist in equilibrium in aqueous solution.
• Hydronium ions and hydroxide ions are the strongest possible acid and base, respectively, that
can exist in aqueous solution.
• Stronger acids react with water to produce hydronium ions and stronger bases react with
water to form hydroxide ions.
• This effect is known as the leveling effect of water.
FORWARD REFERENCES
• Acid-base neutralization reaction will be discussed in detail in Chapter 17 (sections 17.1-
17.3).
16.3 The Autoionization of Water
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• In pure water the following equilibrium is established:
2H2O(l) H3O+(aq) + OH–(aq).
• This process is called the autoionization of water.
The Ion Product of Water
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• We can write an equilibrium constant expression for the autoionization of water.
• Because H2O(l) is a pure liquid, we exclude it from the expression:
Kc = [H3O+][OH–] = Kw.
Chapter 16
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FORWARD REFERENCES
• H2O participating in acid-base reaction as a H+ donor or acceptor will be mentioned in
Chapter 18 (section 18.3).
16.4 The pH Scale
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• In most solutions, [H+] is quite small.
• We express the [H+] in terms of pH.
pH = −log[H+] = −log[H3O+].
• Note that this is a logarithmic scale.
pOH and Other “p” Scales
• We can use a similar system to describe the [OH–].
pOH = −log[OH–].
• Recall that the value of Kw at 25 °C is 1.0 10–14.
• Thus, we can describe a relationship between pH and pOH:
−log[H+]+ (−log[OH–]) = pH + pOH = −logKw = 14.00.
Measuring pH
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• The most accurate method to measure pH is to use a pH meter.
• A pH meter consists of a pair of electrodes connected to a meter that measures small voltages.
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“One-Hundred Years of pH” from Further Readings
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“Do pH in Your Head” from Further Readings
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“pH Estimation” Activity from Instructor’s Resource CD/DVD
Acid-Base Equilibria
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• The relative concentration of the two different forms is sensitive to the pH of the solution.
• Thus, if we know the pH at which the indicator turns color, we can use this color change to
determine whether a solution has a higher or lower pH than this value.
• Some natural products can be used as indicators. (Tea is colorless in acid and brown in base; red
cabbage extract is another natural indicator.)
FORWARD REFERENCES
• Buffer capacity and pH range will be discussed in Chapter 17 (section 17.2).
16.5 Strong Acids and Bases
Strong Acids
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• The most common strong acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4.
• Strong acids are strong electrolytes.
• All strong acids ionize completely in solution.
• For example: nitric acid completely ionizes in water:
HNO3(aq) + H2O(l) → H3O+(aq) + NO3–(aq).
Strong Bases
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• The most common strong bases are ionic hydroxides of the alkali metals or the heavier alkaline earth
metals (e.g., NaOH, KOH, and Sr(OH)2 are all strong bases).
• Strong bases are strong electrolytes and dissociate completely in solution.
FORWARD REFERENCES
• Adding strong acids and bases to buffers will be discussed in detail in Chapter 17 (section
17.2).
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“Colorful Effects of Hydrochloric Acid Dilution” from Live Demonstrations
Chapter 16
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16.6 Weak Acids
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• Weak acids are only partially ionized in aqueous solution.
• There is a mixture of ions and un-ionized acid in solution.
• Therefore, weak acids are in equilibrium:
HA(aq) + H2O(l) H3O+(aq) + A–(aq)
or HA(aq) H+(aq) + A–(aq)
Calculating Ka from pH
• In order to find the value of Ka, we need to know all of the equilibrium concentrations.
• The pH gives the equilibrium concentration of H+.
• Thus, to find Ka we use the pH to find the equilibrium concentration of H+ and then use the
stoichiometric coefficients of the balanced equation to help us determine the equilibrium
concentration of the other species.
• We then substitute these equilibrium concentrations into the equilibrium constant
expression and solve for Ka.
Percent Ionization
• Another measure of acid strength is percent ionization.
Acid-Base Equilibria
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• The higher the percent ionization, the stronger the acid.
• However, we need to keep in mind that percent ionization of a weak acid decreases as the
molarity of the solution increases.
Using Ka to Calculate pH
• Using Ka, we can calculate the concentration of H+ (and hence the pH).
• Write the balanced chemical equation clearly showing the equilibrium.
• Write the equilibrium expression. Look up the value for Ka (in a table).
• Write down the initial and equilibrium concentrations for everything except pure water.
• We usually assume that the equilibrium concentration of H+ is x.
• Substitute into the equilibrium constant expression and solve.
• Remember to convert x to pH if necessary.
• What do we do if we are faced with having to solve a quadratic equation in order to determine the
value of x?
Polyprotic Acids
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• Polyprotic acids have more than one ionizable proton.
• The protons are removed in successive steps.
• Consider the weak acid, H2SO3 (sulfurous acid):
FORWARD REFERENCES
• The use of weak acids and bases as analytes in acid-base titrations will be discussed in
Chapter 17 (section 17.3).
• Absorption of CO2 by the oceans, and subsequent equilibria will be discussed in Chapter 18
(section 18.3).
Chapter 16
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• The role of a triprotic H3PO4 in the formation of nucleic acids will be demonstrated in
Chapter 24 (section 24.10).
16.7 Weak Bases
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• Weak bases remove protons from substances.
• There is an equilibrium between the base and the resulting ions:
weak base + H2O(l) conjugate acid + OH–(aq)
Types of Weak Bases
• Weak bases generally fall into one of two categories.
• Neutral substances with a lone pair of electrons that can accept protons.
FORWARD REFERENCES
• NH3 as a weak base will be further mentioned in Chapter 22 (section 22.5).
• The fact that many organic compounds contain basic groups such as –NH2, –NHR, and –NR2
will be brought up in Chapter 24 (section 24.1).
• Amines, as weak bases, will be discussed in Chapter 24 (section 24.4).
16.8 Relationship Between Ka and Kb
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• We can quantify the relationship between the strength of an acid and the strength of its conjugate
base.
• Consider the following equilibria:
Acid-Base Equilibria
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• The net reaction is the autoionization of water.
H2O(l) H+(aq) + OH–(aq)
• Recall that: Kw = [H+][OH–]
• We can use this information to write an expression that relates the values of Ka, Kb, and Kw for a
conjugate acid-base pair.
16.9 Acid-Base Properties of Salt Solutions
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• Nearly all salts are strong electrolytes.
• Therefore, salts in solution exist entirely of ions.
• Acid-base properties of salts are a consequence of the reactions of their ions in solution.
• Many salt ions can react with water to form OH– or H+.
• This process is called hydrolysis.
An Anion’s Ability to React with Water
• Consider an anion, X–, as the conjugate base of an acid.
• Anions from weak acids are basic.
A Cation’s Ability to React with Water
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• Polyatomic cations that have one or more ionizable protons are conjugate acids of weak bases.
• They tend to decrease pH.
• Many metal ions can cause a decrease in pH.
Chapter 16
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• The larger the charge on the metal ion, the stronger the interaction between the ion and
the oxygen of its hydrating water molecules.
• This weakens the O–H bonds in the water molecules and facilitates proton transfer from
hydration water molecules to solvent water molecules.
Combined Effect of Cation and Anion in Solution
• The pH of a solution may be qualitatively predicted using the following guidelines:
• Salts derived from a strong acid and a strong base are neutral.
• Examples are NaCl and Ba(NO3)2.
16.10 Acid-Base Behavior and Chemical Structure
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Factors that Affect Acid Strength
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• Consider H–X.
• For this substance to be an acid:
• The H–X bond must be polar with H+ and X–.
• In ionic hydrides, the bond polarity is reversed.
Binary Acids
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• The H–X bond strength is important in determining relative acid strength in any group in the periodic
table.
• The H–X bond strength tends to decrease as the element X increases in size.
• Acid strength increases down a group; base strength decreases down a group.
• H–X bond polarity is important in determining relative acid strength in any period of the periodic
table.
Acid-Base Equilibria
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Oxyacids
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• Many acids contain one or more O–H bonds.
• Acids that contain OH groups (and often additional oxygen atoms) bound to the central atom are
called oxyacids.
• All oxyacids have the general structure Y–O–H.
• The strength of the acid depends on Y and the atoms attached to Y.
• As the electronegativity of Y increases, so does the acidity of the substance.
• Example: HClO4 > HClO3 > HClO2 > HClO
Carboxylic Acids
• There is a large class of acids that contain a –COOH group (a carboxyl group).
• Acids that contain this group are called carboxylic acids.
• Examples are acetic acid, benzoic acid, and formic acid.
FORWARD REFERENCES
• Oxoacids and oxoanions containing halogens will be mentioned in Chapter 22 (section 22.4).
• Carboxylic acids will be discussed in Chapter 24 (section 24.4).
• Amphiprotic behavior of amino acids and zwitterions of amino acids will be discussed in
Chapter 24 (section 24.7).
• The relative strength of acetic vs. pyruvic acids (with an additional carbonyl group) will be
discussed in a sample integrative exercise in Chapter 24 (section 24.10).
Chapter 16
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16.11 Lewis Acids and Bases
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• A Brønsted-Lowry acid is a proton donor.
• Focusing on electrons, a Brønsted-Lowry acid can be considered as an electron pair acceptor.
• Lewis proposed a new definition of acids and bases that emphasizes the shared electron pair.
• For example, consider the reaction:
H2O(l) + CO2(g) → H2CO3(aq)
• Water acts as the electron pair donor and carbon dioxide as the electron pair acceptor in this
reaction.
• Overall, the water (Lewis base) has donated a pair of electrons to the CO2 (Lewis acid).
increases (e.g., Na+ vs. Ca2+ and Zn2+ vs. Al3+).
FORWARD REFERENCES
• Lewis bases and Crystal Field Theory will be discussed in Chapter 23 (section 23.6).
Acid-Base Equilibria
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Further Readings:
1. Doris Kolb, “Acids and Bases,” J. Chem. Educ., Vol. 55, 1978, 459–464.
2. William B. Jensen, “The Origin of the Term ‘Base’,” J. Chem. Educ., Vol. 83, 2006, 1130.
5. Jamie L. Adcock, "Teaching Brønsted-Lowry Acid–Base Theory in a Direct Comprehensive Way,” J.
Chem. Educ., Vol. 78, 2001, 1495–1496.
6. A. M. de Lange and J. H. Potgieter, “Acid and Base Dissociation Constants of Water and Its
Associated Ions,” J. Chem. Educ., Vol. 68, 1991, 304–305.
11. Stephen J. Hawkes, “Teaching the Truth about pH,” J. Chem. Educ., Vol. 71, 1994, 747–749.
12. Doris Kolb, “The pH Concept,” J. Chem. Educ., Vol. 56, 1979, 49–53.
13. Richard F. Burton, “Defining and Teaching pH,” J. Chem. Educ., Vol. 84, 2007, 1129.
17. Todd P. Silverstein, “Weak vs. Strong Acids and Bases: The Football Analogy,” J. Chem. Educ., Vol.
77, 2000, 849–850.
Chapter 16
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Live Demonstrations:
1. Bassam Z. Shakhashiri, “Acid-Base Indicators Extracted from Plants,” Chemical Demonstrations: A
Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989), pp.
50–57. A wide range of plant materials is used as sources of acid-base indicators. Included in this group
is an old favorite: red cabbage.
4. Bassam Z. Shakhashiri, “Colorful Acid-Base Indicators,” Chemical Demonstrations: A Handbook for
Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989), pp. 33–40.
5. Bassam Z. Shakhashiri, “Hydrolysis: Acidic and Basic Properties of Salts,” Chemical Demonstrations:
A Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989),
pp. 103–108.
8. Bassam Z. Shakhashiri, “Rainbow Colors with Mixed Acid-Base Indicators,” Chemical
Demonstrations: A Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of
Wisconsin Press, 1989), pp. 41–46.
