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5: Reactions of Metals with Nonmetals (Oxidation–Reduction)
C H E M I S T R Y I N F OCUS
Garden-Variety Acid–Base Indicators
When placed in an acidic solution, most of
the basic form of the indicator is converted to
the acidic form by the reaction
can flowers tell us about acids and
bases? Actually, some flowers can tell us whether
the soil they are growing in is acidic or basic. For
example, in acidic soil, bigleaf hydrangea blossoms will be blue; in basic (alkaline) soil, the
flowers will be red. What is the secret? The pigment in the flower is an acid–base indicator.
Generally, acid–base indicators are dyes that
are weak acids. Because indicators are usually
complex molecules, we often symbolize them as
HIn. The reaction of the indicator with water can
be written as
HIn(aq) ϩ H2O(l) 3
4 H3Oϩ(aq) ϩ InϪ(aq)
To work as an acid–base indicator, the conjugate acid–base forms of these dyes must have different colors. The acidity level of the solution will
determine whether the indicator is present mainly
in its acidic form (HIn) or its basic form (InϪ).
InϪ(aq) ϩ Hϩ(aq) n HIn(aq)
When placed in a basic solution, most of the
acidic form of the indicator is converted to the
basic form by the reaction
HIn(aq) ϩ OHϪ(aq) n InϪ(aq) ϩ H2O(l)
It turns out that many fruits, vegetables,
and flowers can act as acid–base indicators. Red,
blue, and purple plants often contain a class of
chemicals called anthocyanins, which change
color based on the acidity level of the surroundings. Perhaps the most famous of these plants is
red cabbage. Red cabbage contains a mixture of
anthocyanins and other pigments that allow it
to be used as a “universal indicator.” Red cabbage juice appears deep red at a pH of 1–2, purple at a pH of 4, blue at a pH of 8, and green at
a pH of 11.
16.5 Calculating the pH of Strong
To learn to calculate the pH of solutions of strong acids.
In this section we will learn to calculate the pH for a solution containing a
strong acid of known concentration. For example, if we know a solution contains 1.0 M HCl, how can we find the pH of the solution? To answer this
question we must know that when HCl dissolves in water, each molecule dissociates (ionizes) into Hϩ and ClϪ ions. That is, we must know that HCl is a
strong acid. Thus, although the label on the bottle says 1.0 M HCl, the solution contains virtually no HCl molecules. A 1.0 M HCl solution contains Hϩ
and ClϪ ions rather than HCl molecules. Typically, container labels indicate
the substance(s) used to make up the solution but do not necessarily describe
the solution components after dissolution. In this case,
1.0 M HCl S 1.0 M Hϩ and 1.0 M ClϪ
Therefore, the [Hϩ] in the solution is 1.0 M. The pH is then
pH ϭ Ϫlog[Hϩ] ϭ Ϫlog(1.0) ϭ 0
variety of flower petals, including delphiniums,
geraniums, morning glories, and, of course, hydrangeas.
Neil Holmes/Homes Garden Photos/Alamy
Other natural indicators include the skins
of beets (which change from red to purple in
very basic solutions), blueberries (which change
from blue to red in acidic solutions), and a wide
Calculating the pH of Strong Acid Solutions
Calculate the pH of 0.10 M HNO3.
HNO3 is a strong acid, so the ions in solution are Hϩ and NO3Ϫ. In this case,
0.10 M HNO3 S 0.10 M Hϩ and 0.10 M NOϪ
[Hϩ] ϭ 0.10 M
pH ϭ Ϫlog(0.10) ϭ 1.00
Self-Check EXERCISE 16.7 Calculate the pH of a solution of 5.0 ϫ 10Ϫ3 M HCl.
See Problems 16.57 and 16.58. ■
534 Chapter 16 Acids and Bases
16.6 Buffered Solutions
Water: pH ϭ 7
0.01 M HCl: pH ϭ 2
To understand the general characteristics of buffered solutions.
A buffered solution is one that resists a change in its pH even when a
strong acid or base is added to it. For example, when 0.01 mole of HCl is
added to 1 L of pure water, the pH changes from its initial value of 7 to 2, a
change of 5 pH units. However, when 0.01 mole of HCl is added to a solution containing both 0.1 M acetic acid (HC2H3O2) and 0.1 M sodium acetate
(NaC2H3O2), the pH changes from an initial value of 4.74 to 4.66, a change
of only 0.08 pH unit. The latter solution is buffered—it undergoes only a
very slight change in pH when a strong acid or base is added to it.
Buffered solutions are vitally important to living organisms whose cells
can survive only in a very narrow pH range. Many goldfish have died because their owners did not realize the importance of buffering the aquarium
water at an appropriate pH. For humans to survive, the pH of the blood must
be maintained between 7.35 and 7.45. This narrow range is maintained by
several different buffering systems.
A solution is buffered by the presence of a weak acid and its conjugate
base. An example of a buffered solution is an aqueous solution that contains
acetic acid and sodium acetate. The sodium acetate is a salt that furnishes acetate ions (the conjugate base of acetic acid) when it dissolves. To see how
this system acts as a buffer, we must recognize that the species present in this
Hans Reinhard/Bruce Coleman, Inc.
When NaC2H3O2 is
dissolved, it produces
the separated ions
What happens in this solution when a strong acid such as HCl is added?
In pure water, the Hϩ ions from the HCl would accumulate, thus lowering
HCl ¡ Hϩ ϩ ClϪ
For goldfish to survive, the pH
of the water must be carefully
However, this buffered solution contains C2H3O2Ϫ ions, which are basic.
That is, C2H3O2Ϫ has a strong affinity for Hϩ, as evidenced by the fact that
HC2H3O2 is a weak acid. This means that the C2H3O2Ϫ and Hϩ ions do not
exist together in large numbers. Because the C2H3O2Ϫ ion has a high affinity
for Hϩ, these two combine to form HC2H3O2 molecules. Thus the Hϩ from
the added HCl does not accumulate in solution but reacts with the C2H3O2Ϫ
H ϩ (aq) ϩ C2H3OϪ
2 (aq) S HC2H3O2(aq)
Next consider what happens when a strong base such as sodium hydroxide is added to the buffered solution. If this base were added to pure water, the OHϪ ions from the solid would accumulate and greatly change (raise)
NaOH ¡ Naϩ ϩ OHϪ
However, in the buffered solution the OHϪ ion, which has a very strong affinity for Hϩ, reacts with HC2H3O2 molecules as follows:
HC2H3O2(aq) ϩ OHϪ(aq) S H2O(l) ϩ C2H3OϪ
Table 16.3 The Characteristics of a Buffer
1. The solution contains a weak acid HA and its conjugate base AϪ.
2. The buffer resists changes in pH by reacting with any added Hϩ or OHϪ so
that these ions do not accumulate.
3. Any added Hϩ reacts with the base AϪ.
Hϩ(aq) ϩ AϪ(aq) n HA(aq)
4. Any added OHϪ reacts with the weak acid HA.
OHϪ(aq) ϩ HA(aq) n H2O(l) ϩ AϪ(aq)
This happens because, although C2H3O2Ϫ has a strong affinity for Hϩ, OHϪ
has a much stronger affinity for Hϩ and thus can remove Hϩ ions from acetic
Note that the buffering materials dissolved in the solution prevent
added Hϩ or OHϪ from building up in the solution. Any added Hϩ is trapped
by C2H3O2Ϫ to form HC2H3O2. Any added OHϪ reacts with HC2H3O2 to form
H2O and C2H3O2Ϫ.
The general properties of a buffered solution are summarized in
C H A P T E R
Arrhenius concept of
acids and bases (16.1)
conjugate acid (16.1)
conjugate base (16.1)
hydronium ion (16.1)
strong acid (16.2)
weak acid (16.2)
diprotic acid (16.2)
organic acid (16.2)
carboxyl group (16.2)
ionization of water (16.3)
neutral solution (16.3)
acidic solution (16.3)
basic solution (16.3)
pH scale (16.4)
buffered solution (16.6)
1. Acids or bases in water are commonly described by
two different models. Arrhenius postulated that acids
produce Hϩ ions in aqueous solutions and that bases
produce OHϪ ions. The Brønsted–Lowry model is
directs you to the Chemistry in Focus feature in the chapter
indicates visual problems
interactive versions of these problems are assignable in OWL
more general: an acid is a proton donor, and a base is
a proton acceptor. Water acts as a Brønsted–Lowry
base when it accepts a proton from an acid to form a
HA(aq) ϩ H2O(l) 3
4 H3Oϩ(aq) ϩ AϪ(aq)
A conjugate base is everything that remains of the
acid molecule after the proton is lost. A conjugate
acid is formed when a proton is transferred to the
base. Two substances related in this way are called a
conjugate acid–base pair.
2. A strong acid or base is one that is completely ionized
(dissociated). A weak acid is one that is ionized (dissociated) only to a slight extent. Strong acids have
weak conjugate bases. Weak acids have relatively
strong conjugate bases.
3. Water is an amphoteric substance—it can behave either as an acid or as a base. The ionization of water reveals this property; one water molecule transfers a
536 Chapter 16 Acids and Bases
proton to another water molecule to produce a hydronium ion and a hydroxide ion.
H2O(l) ϩ H2O(l) 3
4 H3Oϩ(aq) ϩ OHϪ(aq)
Kw ϭ [H3O ][OH ] ϭ [H ][OH ]
is called the ion-product constant. It has been shown
experimentally that at 25 °C,
[Hϩ] ϭ [OHϪ] ϭ 1.0 ϫ 10Ϫ7 M
so Kw ϭ 1.0 ϫ 10Ϫ14.
4. In an acidic solution, [Hϩ] is greater than [OHϪ]. In a
basic solution, [OHϪ] is greater than [Hϩ]. In a neutral
solution, [Hϩ] ϭ [OHϪ].
5. To describe [H ] in aqueous solutions, we use the pH
pH ϭ Ϫlog[Hϩ]
Note that the pH decreases as [Hϩ] (acidity) increases.
6. The pH of strong acid solutions can be calculated directly from the concentration of the acid, because
100% dissociation occurs in aqueous solution.
7. A buffered solution is one that resists a change in its
pH even when a strong acid or base is added to it. A
buffered solution contains a weak acid and its conjugate base.
Active Learning Questions
These questions are designed to be considered by groups of
students in class. Often these questions work well for introducing a particular topic in class.
1. You are asked for the Hϩ concentration in a solution
of NaOH(aq). Because sodium hydroxide is a strong
base, can we say there is no Hϩ, since having Hϩ
would imply that the solution is acidic?
2. Explain why ClϪ does not affect the pH of an aqueous
8. Can the pH of a solution be negative? Explain.
9. Stanley’s grade-point average (GPA) is 3.28. What is
10. A friend asks the following: “Consider a buffered solution made up of the weak acid HA and its salt NaA.
If a strong base like NaOH is added, the HA reacts
with the OHϪ to make AϪ. Thus, the amount of acid
(HA) is decreased, and the amount of base (AϪ) is increased. Analogously, adding HCl to the buffered solution forms more of the acid (HA) by reacting with
the base (AϪ). How can we claim that a buffered solution resists changes in the pH of the solution?” How
would you explain buffering to your friend?
11. Mixing together aqueous solutions of acetic acid and
sodium hydroxide can make a buffered solution.
12. Could a buffered solution be made by mixing aqueous solutions of HCl and NaOH? Explain.
13. Consider the equation: HA(aq) ϩ H2O 3
4 H3Oϩ(aq) ϩ
a. If water is a better base than AϪ, which way will
b. If water is a better base than AϪ, does this mean
that HA is a strong or a weak acid?
c. If water is a better base than AϪ, is the value for Ka
greater or less than 1?
14. Choose the answer that best completes the following
statement and defend your answer. When 100.0 mL
of water is added to 100.0 mL of 1.00 M HCl,
a. the pH decreases because the solution is diluted.
b. the pH does not change because water is neutral.
c. the pH is doubled because the volume is now
d. the pH increases because the concentration of Hϩ
e. the solution is completely neutralized.
15. You mix a solution of a strong acid with a pH of 4 and
an equal volume of a strong acid solution with a pH
of 6. Is the final pH less than 4, between 4 and 5, 5,
between 5 and 6, or greater than 6? Explain.
3. Write the general reaction for an acid acting in water.
What is the base in this case? The conjugate acid? The VP 16. The following figures are molecular-level representaconjugate base?
tions of acid solutions. Label each as a strong acid or
a weak acid.
4. Differentiate among the terms concentrated, dilute,
weak, and strong in describing acids. Use molecularlevel pictures to support your answer.
5. What is meant by “pH”? True or false: A strong acid
always has a lower pH than a weak acid does. Explain.
6. Consider two separate solutions: one containing a
weak acid, HA, and one containing HCl. Assume that
you start with 10 molecules of each.
a. Draw a molecular-level picture of what each solution looks like.
b. Arrange the following from strongest to weakest
base: ClϪ, H2O, AϪ. Explain.
7. Why is the pH of water at 25 °C equal to 7.00?
17. Answer the following questions concerning buffered
a. Explain what a buffered solution does.
b. Describe the substances that make up a buffered
c. Explain how a buffered solution works.