Chapter15

Contents and Concepts
Acid–Base Concepts
1.Arrhenius Concept of Acids and Bases
2.Brønsted–Lowry Concept of Acids and Bases
3.Lewis Concept of Acids and Bases
Acid and Base Strengths
4.Relative Strengths of Acids and Bases
5.Molecular Structure and Acid Strength
Copyright © Cengage Learning. All rights reserved.

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Self-Ionization of Water and pH
6.Self-Ionization of Water
7.Solutions of a Strong Acid or Base
8.The pH of a Solution


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Learning Objectives
Acid Base Concepts
Arrhenius Concept of Acids and Base
a. Define acid and base according to the
Arrhenius concept.
Brønsted–Lowry Concept of Acids and Bases
a. Define acid and base according to the
Brønsted–Lowry concept.
b. Define the term conjugate acid–base pair.
c. Identify acid and base species.
d. Define amphiprotic species.

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3. Lewis Concept of Acids and Bases
a. Define Lewis acid and Lewis base.
b. Identify Lewis acid and Lewis base
species.
Acid and Base Strengths
4. Relative Strengths of Acids and Bases
a. Understand the relationship between the
strength of an acid and that of its conjugate
base.
b. Decide whether reactants or products are
favored in an acid–base reaction.

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5. Molecular Structure and Acid Strength
a. Note the two factors that determine relative
acid strengths.
b. Understand the periodic trends in the
strengths of the binary acids HX.
c. Understand the rules for determining the
relative strengths of oxoacids.
d. Understand the relative acid strengths of a
polyprotic acid and its anions.


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Self-Ionization of Water and pH
6. Self-Ionization of Water
a. Define self-ionization (or autoionization).
b. Define the ion-product constant for water.
7. Solutions of a Strong Acid or Base
a. Calculate the concentrations of H3O+ and
OH- in solutions of a strong acid or base.


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8. The pH of a Solution
a. Define pH.
b. Calculate the pH from the hydronium-ion
concentration.
c. Calculate the hydronium-ion concentration
from the pH.
d. Describe the determination of pH by a pH
meter and by acid–base indicators.


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When gaseous hydrogen chloride meets gaseous
ammonia, a smoke composed of ammonium
chloride is formed.
HCl(g) + NH3(g) → NH4Cl(s)
SHOW DIFFUSION VIDEO NOW!!!
This is an acid–base reaction.


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We will examine three ways to explain acid–base
behavior:
H+ and OH−

Arrhenius Concept
donor
H+ = proton acceptor
Brønsted–Lowry Concept
Lewis Concept electron pair donor
acceptor
acid
base
Note: H+ in water is H O+
3


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Arrhenius Concept of Acids and Bases
An Arrhenius acid is a substance that, when
dissolved in water, increases the concentration of
hydronium ion, H3O+(aq).
An Arrhenius base is a substance that, when
dissolved in water, increases the concentration of
hydroxide ion, OH-(aq).


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The Arrhenius concept limits bases to compounds
that contain a hydroxide ion.
The Brønsted–Lowry concept expands the
compounds that can be considered acids and
bases.


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Brønsted–Lowry Concept of Acids and Bases
An acid–base reaction is considered a proton (H+)
transfer reaction.
H+
H+

H+
H+


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A Brønsted–Lowry acid is the species donating a
proton in a proton-transfer reaction; it is a proton
donor.
A Brønsted–Lowry base is the species accepting a
proton in a proton-transfer reaction; it is a proton
acceptor.

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Substances in the acid–base reaction that differ by
the gain or loss of a proton, H+, are called a
conjugate acid–base pair. The acid is called the
conjugate acid; the base is called a conjugate
base.

Acid

Base


Conjugate Conjugate
acid
base
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?

What is the conjugate acid of H2O?
What is the conjugate base of H2O?

The conjugate acid of H2O has gained a proton.
It is H3O+.
The conjugate base of H2O has lost a proton.
It is OH-.

Label each species as an acid or base. Identify the
conjugate acid-base pairs.

a.

HCO3−(aq) + HF(aq)
Base

b.

Acid

HCO3−(aq) + OH−(aq)

Acid


Base

H2CO3(aq) + F−(aq)

Conjugate Conjugate
acid
base

CO32−(aq) + H2O(l)
Conjugate Conjugate
base
acid
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Species that can act as both an acid and a base
are called amphiprotic or amphoteric species.
Identify any amphiprotic species in the previous
problem.
HCO3− was a base in the first reaction and an acid
in the second reaction. It is amphiprotic.


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Lewis Concept of Acids and Bases
A Lewis acid is a species that can form a covalent
bond by accepting an electron pair from another
species.
A Lewis base is a species that can form a covalent
bond by donating an electron pair to another
species.


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Relative Strengths of Acids and Bases
The stronger an acid, the weaker its conjugate
base.
The weaker an acid, the stronger its conjugate
base.


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?

Formic acid, HCHO2, is a stronger acid
than acetic acid, HC2H3O2. Which is the
stronger base: formate ion, CHO2−, or
acetate ion, C2H3O2−?

Because formic acid is stronger than acetic acid,
formate ion (which is the conjugate base of formic
acid) will be a weaker base than acetate ion (which
is the conjugate base of acetic acid).
The acetate ion is a stronger base than the
formate ion.

Molecular Structure and Acid Strength
The strength of an acid depends on how easily the
proton, H+, is lost or removed. The more polarized
the bond between H and the atom to which it is
bonded, the more easily the H+ is lost or donated.
We will look now at factors that affect how easily
the hydrogen can be lost and, therefore, acid
strength.


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For a binary acid, as the size of X in HX increases,
going down a group, acid strength increases.


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For a binary acid, going across a period, as the
electronegativity increases, acid strength
increases.


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?

Which is a stronger acid: HF or HCl?
Which is a stronger acid: H2O or H2S?
Which is a stronger acid: HCl or H2S?

HF and HCl
These are binary acids from the same group,
so we compare the size of F and Cl. Because
Cl is larger, HCl is the stronger acid.

H2O and H2S
These are binary acids from the same group, so
we compare the size of O and S. Because S is
larger, H2S is the stronger acid.
HCl and H2S
These are binary acids from the same period, but
different groups, so we compare the
electronegativity of Cl and S. Because Cl is more
electronegative, HCl is the stronger acid.

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For oxoacids, several factors are relevant: the
number and bonding of oxygens, the central
element, and the charge on the species.
For a series of oxoacids, (OH)mYOn, acid strength
increases as n increases.
(OH)ClO
(OH)ClO2
(OH)ClO3
(OH)Cl
n=0

n=1

Weakest

n=2

n=3

Strongest
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For a series of oxoacids differing only in the central
atom Y, the acid strength increases with the
electronegativity of Y.

Stronger

Weaker
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The acid strength of a polyprotic acid and its
anions decreases with increasing negative charge.
H2CO3 is a stronger acid than HCO3−.
H2SO4 is a stronger acid than HSO4−.
H3PO4 is a stronger acid than H2PO4−.
H2PO4- is a stronger acid than HPO42−.


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A reaction will always go in the direction from
stronger acid to weaker acid, and from stronger
base to weaker base.


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?

Decide which species are favored at
the completion of the following
reaction:
HCN(aq) + HSO3−(aq) 
CN−(aq) + H2SO3(aq)

We first identify the acid on each side of the
reaction: HCN and H2SO3.
Next, we compare their acid strength: H2SO3 is
stronger.
This reaction will go from right to left (←), and the
reactants are favored.

Self-Ionization of Water

H2O(l) + H2O(l)
Base

Acid


H3O+(aq) + OH-(aq)
Conjugate
acid

Conjugate
base

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H2O(l) + H2O(l)  H3O+(aq) + OH−(aq)
We call the equilibrium constant the ion-product
constant, Kw.
Kw = [H3O+][OH−]
At 25°C, Kw = 1.0 × 10−14
As temperature increases, the value of Kw
increases.

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Solutions of a Strong Acid or Strong Base
The concentration of hydronium or hydroxide in a
solution of strong acid or base is related to the
stoichiometry of the acid or base.


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Calculate the hydronium and hydroxide
ion concentration at 25°C in
a. 0.10 M HCl
b. 1.4 × 10−4 M Mg(OH)2
a. When HCl ionizes, it gives H+ and Cl−.
So [H+] = [Cl−] = [HCl] = 0.10 M.
b. When Mg(OH)2 ionizes, it gives Mg2+ and 2 OH−.
So [OH−] = 2[Mg2+] = 2[Mg(OH)2] = 2.8 × 10−4 M.

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Solutions can be characterized as
Acidic:
[H3O+] > 1.0 × 10−7 M
Neutral:

[H3O+] = 1.0 × 10−7 M

Basic:

[H3O+] < 1.0 × 10−7 M


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Concept Check 15.3
Rank the following solutions from most acidic to
most basic (water molecules have been omitted for
clarity).

A has 5 H3O+ and 5 OH-. It is neutral.
B has 7 H3O+ and 3 OH-. It is acidic.
C has 3 H3O+ and 7 OH-. It is basic.
Listed from most acidic to most basic: B, A, C.

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The pH of a Solution
pH = –log[H3O+]
For a log, only the decimal part of the number has
significant figures. The whole number part, called
the characteristic, is not significant.


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?

Calculate the pH of typical adult blood,
which has a hydronium ion
concentration of 4.0 × 10−8 M.
[H3O+] = 4.0 × 10−8 M
pH = –log [H3O+]
pH = – log (4.0 × 10−8) = – (– 7.40)
pH = 7.40

Note: The two significant figures are the two
decimal places.

pOH = –log[OH−]
pH + pOH = 14.00 (at 25°C)


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[H3O+] = 10−pH
[OH−] = 10−pOH


15 | 44

The pH of natural rain in 5.60. What is
its hydronium ion concentration?
pH = 5.60
[H3O+] = 10−pH = 10−5.60
[H3O+] = 2.5 × 10-6 M
Because the pH has two decimal places, the
concentration can have only two significant figures.

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The next slide summarizes the conversions
involving H3O+, OH−, pH, and pOH.
Note that you can only go around the edges of the
square; it takes two steps to go from one corner to
the opposite corner.


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1× 10 −14
[OH ] =
[H3O + ]
−

[H3O ]
+

[H3O + ] =
10

−pH

[OH ]
−

1× 10 −14
[H3O ] =
[OH− ]
+

pH =

pOH =

[OH−] =
+

− log[H3O ]

pH

10

pOH = 14 − pH

−pOH

−log[OH− ]

pOH

pH = 14 − pOH

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