Acids and Bases
Section 18.1 Introduction to Acids
and Bases
Section 18.2 Strengths of Acids
and Bases
Section 18.3 Hydrogen Ions
and pH
Section 18.4 Neutralization
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Section 18.1 Introduction to Acids and Bases
• Identify the physical and chemical properties of acids
and bases.
• Classify solutions as acidic, basic, or neutral.
• Compare the Arrhenius, Brønsted-Lowry, and Lewis
models of acids and bases.
Lewis structure: a model that uses electron-dot
structures to show how electrons are arranged in
molecules
Section 18.1 Introduction to Acids and Bases
(cont.)
acidic solution
conjugate base
basic solution
conjugate acid-base pair
Arrhenius model
amphoteric
Brønsted-Lowry model
Lewis model
conjugate acid
Different models help describe the
behavior of acids and bases.
Properties of Acids and Bases
• Acids taste sour. Bases taste bitter and feel
slippery.
• Acids and bases are conductors of electricity.
• Acids and bases can be identified by their
reactions with some metals and metal
carbonates.
Properties of Acids and Bases (cont.)
• Acids turn blue litmus red.
• Bases turn red litmus blue.
• Magnesium and zinc react with acids to
produce hydrogen gas.
• Geologists identify limestone because it
produces bubbles of carbon dioxide when
exposed to hydrochloric acid.
Properties of Acids and Bases (cont.)
• All water solutions contain hydrogen ions
(H+) and hydroxide ions (OH–).
• An acidic solution contains more hydrogen
ions than hydroxide ions.
• A basic solution contains more hydroxide
ions than hydrogen ions.
Properties of Acids and Bases (cont.)
• The usual solvent for acids and bases is
water—water produces equal numbers of
hydrogen and hydroxide ions in a process
called self-ionization.
H2O(l) + H2O(l) ↔ H3O+(aq) + OH–(aq)
• The hydronium ion is H3O+.
The Arrhenius Model
• The Arrhenius model states that an acid
is a substance that contains hydrogen and
ionizes to produce hydrogen ions in
aqueous solution, and a base is a
substance that contains a hydroxide group
and dissociates to produce a hydroxide ion
in solution.
The Arrhenius Model (cont.)
• Arrhenius acids and bases
– HCl ionizes to produce H+ ions.
– HCl(g) → H+(aq) + Cl–(aq)
– NaOH dissociates to produce OH– ions.
– NaOH(s) → Na+(aq) + OH–(aq)
– Some solutions produce hydroxide ions even
though they do not contain a hydroxide group.
The Brønsted-Lowry Model
• The Brønsted-Lowry Model of acids and
bases states that an acid is a hydrogen ion
donor, and a base is a hydrogen ion
acceptor.
• The Brønsted-Lowry Model is a more
inclusive model of acids and bases.
The Brønsted-Lowry Model (cont.)
• A conjugate acid is the species produced
when a base accepts a hydrogen ion.
• A conjugate base is the species produced
when an acid donates a hydrogen ion.
• A conjugate acid-base pair consists of two
substances related to each other by donating
and accepting a single hydrogen ion.
The Brønsted-Lowry Model (cont.)
• Hydrogen fluoride—a Brønsted-Lowry acid
– HF(aq) + H2O(l) ↔ H3O+(aq) + F–(aq)
– HF = acid, H2O = base, H3O+ = conjugate acid,
F– = conjugate base
The Brønsted-Lowry Model (cont.)
• Ammonia— Brønsted-Lowry base
– NH3(aq) + H2O(l) ↔ NH4+(aq) + OH–(aq)
– NH3 = base, H2O(l) = acid, NH4+ = conjugate
acid, OH– = conjugate base
• Water and other substances that can act as
acids or bases are called amphoteric.
Monoprotic and Polyprotic Acids
• An acid that can donate only one hydrogen
ion is a monoprotic acid.
• Only ionizable hydrogen atoms can be
donated.
Monoprotic and Polyprotic Acids (cont.)
• Acids that can donate more than one
hydrogen ion are polyprotic acids.
The Lewis Model
• According to the Lewis model, a Lewis
acid is an electron-pair acceptor and a
Lewis base is an electron pair donor.
• The Lewis model includes all the substances
classified as Brønsted-Lowry acids and bases
and many more.
Section 18.1 Assessment
A Lewis acid is a(n) ____.
A. electron pair donor
B. hydrogen ion donor
C. electron pair acceptor
D
C
A
0%
B
D. substance that contains
an hydroxide group
A. A
B. B
C. C
0%
0%
0%
D. D
Section 18.1 Assessment
A conjugate acid is formed when:
A. a base accepts a hydrogen ion
B. an acid accepts a hydrogen ion
C. an acid donates a hydrogen ion
D
C
A
0%
B
D. a base donates a hydrogen ion
A. A
B. B
C. C
0%
0%
0%
D. D
Section 18.2 Strengths of Acids and Bases
• Relate the strength of an
acid or base to its
degree of ionization.
• Compare the strength of
a weak acid with the
strength of its conjugate
base.
• Explain the relationship
between the strengths of
acids and bases and the
values of their ionization
constants.
electrolyte: an ionic
compound whose
aqueous solution
conducts an electric
current
Section 18.2 Strengths of Acids and Bases (cont.)
strong acid
weak acid
acid ionization constant
strong base
weak base
base ionization constant
In solution, strong acids and bases
ionize completely, but weak acids and
bases ionize only partially.
Strengths of Acids
• Acids that ionize completely are strong
acids.
• Because they produce the maximum number
of hydrogen ions, strong acids are good
conductors of electricity.
Strengths of Acids (cont.)
• Acids that ionize only partially in dilute
aqueous solutions are called weak acids.
Strengths of Acids (cont.)
• With a strong acid, the conjugate base is a
weak base.
• Equilibrium lies almost completely to the right
in the equation because the conjugate base
has a weaker attraction for the H+ ion than
does the base in the forward reaction.
• In a weak acid, the ionization equilibrium lies
to the far left in the ionization equation
because the conjugate base has a greater
attraction for H+ ions than does the base in
the forward reaction.
Strengths of Acids (cont.)
• The equilibrium constant, Keq, provides a
quantitative measure of the degree of
ionization of an acid.
• The acid ionization constant is the value of
the equilibrium constant expression for the
ionization of a weak acid, Ka.
• Ka indicates whether products or reactants
are favored at equilibrium.
Strengths of Acids (cont.)
• For weak acids, the products tend to be
smaller compared to the un-ionized
molecules (reactant).
• Weaker acids have a smaller Ka.
Strengths of Bases
• A base that dissociates
completely into metal
ions and hydroxide ions
is known as a strong
base.
• A weak base ionizes
only partially in dilute
aqueous solution.
Strengths of Bases (cont.)
• The base ionization constant, Kb, is the
value of the equilibrium constant
expression for the ionization of a base.
Section 18.2 Assessment
A solution with a small Kb is a ____.
A. weak acid
B. weak base
C. strong acid
D
C
A
0%
B
D. strong base
A. A
B. B
C. C
0%
0%
0%
D. D
Section 18.2 Assessment
Where is the equilibrium point in the
ionization equation for a strong acid?
A. far right
B. far left
D
A
0%
C
D. slightly left
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. slightly right
Section 18.3 Hydrogen Ions and pH
• Explain pH and pOH.
• Relate pH and pOH to the
ion product constant for
water.
• Calculate the pH and
pOH of aqueous solutions.
Le Châtelier’s principle:
states that if a stress is
applied to a system at
equilibrium, the system
shifts in the direction that
relieves the stress
ion product constant
pH and pOH are logarithmic for water
scales that express the
pH
concentrations of hydrogen
pOH
ions and hydroxide ions in
aqueous solutions.
Ion Product Constant for Water
• Pure water contains equal concentrations
of H+ and OH– ions.
• The ion production of water, Kw = [H+][OH–].
• The ion product constant for water is the
value of the equilibrium constant expression
for the self-ionization of water.
Ion Product Constant for Water (cont.)
• With pure water at 398 K, both [H+] and
[OH–] are equal to 1.0 × 10–7M.
Kw at 298 K = 1.0 × 10–14
• Kw and LeChâtelier’s Principle proves
[H+] × [OH–] must equal 1.0 × 10–14 at 298 K,
and as [H+] goes up, [OH–] must go down.
pH and pOH
• Concentrations of H+ ions are often
small numbers expressed in
exponential notation.
• pH is the negative logarithm of the
hydrogen ion concentration of a
solution.
pH = –log [H+]
pH and pOH (cont.)
• pOH of a solution is the negative logarithm
of the hydroxide ion concentration.
• pOH = –log [OH–]
• The sum of pH and pOH equals 14.
pH and pOH (cont.)
• For all strong monoprotic acids, the
concentration of the acid is the
concentration of H+ ions.
• For all strong bases, the concentration of
the OH– ions available is the concentration
of OH–.
• Weak acids and weak bases only partially
ionize and Ka and Kb values must be used.
pH and pOH (cont.)
• Litmus paper and a pH meter with
electrodes can determine the pH of a
solution.
Section 18.3 Assessment
In dilute aqueous solution, as [H+]
increases:
A. pH decreases
B. pOH increases
D
A
0%
C
D. all of the above
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. [OH–] decreases
Section 18.3 Assessment
What is the pH of a neutral solution such
as pure water?
A. 0
B. 7
D
A
0%
C
D. 1.0 × 10–14
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 14
Section 18.4 Neutralization
• Write chemical
equations for
neutralization reactions.
• Explain how
neutralization reactions
are used in acid-base
titrations.
• Compare the properties
of buffered and
unbuffered solutions.
stoichiometry: the
study of quantitative
relationships between
the amounts of
reactants used and
products formed by a
chemical reaction; is
based on the law of
conservation of mass
Section 18.4 Neutralization (cont.)
neutralization reaction
acid-base indicator
salt
end point
titration
salt hydrolysis
titrant
buffer
equivalence point
buffer capacity
In a neutralization reaction, an acid
reacts with a base to produce a salt
and water.
Reactions Between Acids and Bases
• A neutralization reaction is a reaction in
which an acid and a base in an aqueous
solution react to produce a salt and water.
• A salt is an ionic compound made up of a
cation from a base and an anion from an
acid.
• Neutralization is a double-replacement
reaction.
Reactions Between Acids and Bases (cont.)
Reactions Between Acids and Bases (cont.)
• Titration is a method for determining the
concentration of a solution by reacting a
known volume of that solution with a
solution of known concentration.
Reactions Between Acids and Bases (cont.)
• In a titration procedure, a measured
volume of an acid or base of unknown
concentration is placed in a beaker, and
initial pH recorded.
• A buret is filled with the titrating solution of
known concentration, called a titrant.
Reactions Between Acids and Bases (cont.)
• Measured volumes of the standard solution
are added slowly and mixed into the solution
in the beaker, and the pH is read and
recorded after each addition. The process
continues until the reaction reaches the
equivalence point, which is the point at
which moles of H+ ion from the acid equals
moles of OH– ion from the base.
• An abrupt change in pH occurs at the
equivalence point.
Reactions Between Acids and Bases (cont.)
Reactions Between Acids and Bases (cont.)
• Chemical dyes whose color are affected by
acidic and basic solutions are called acidbase indicators.
Reactions Between Acids and Bases (cont.)
• An end point is the point at which an
indicator used in a titration changes color.
• An indicator will change color at the
equivalence point.
Salt Hydrolysis
• In salt hydrolysis, the anions of the
dissociated salt accept hydrogen ions from
water or the cations of the dissociated salt
donate hydrogen ions to water.
Salt Hydrolysis
(cont.)
• Salts that produce basic solutions
− KF is the salt of a strong base (KOH) and a weak
acid (HF).
KF(s) → K+(aq) + F–(aq)
Salt Hydrolysis
(cont.)
• Salts that produce acidic solutions
− NH4Cl is the salt of a weak base (NH3) and strong
acid (HCl).
− When dissolved in water, the salt dissociates into
ammonium ions and chloride ions.
NH4Cl(s) → NH4+(aq) + Cl–(aq)
Salt Hydrolysis
(cont.)
• Salts that produce neutral solutions
− NaNO3 is the salt of a strong acid (HNO3) and a
strong base (NaOH).
− Little or no salt hydrolysis occurs because neither
Na+ nor NO3– react with water.
Buffered Solutions
• The pH of blood must be kept in within a
narrow range.
• Buffers are solutions that resist changes in
pH when limited amounts of acid or base are
added.
Buffered Solutions (cont.)
• Ions and molecules in a buffer solution
resist changes in pH by reacting with any
hydrogen ions of hydroxide ions added to
the buffered solution.
HF(aq) ↔ H+(aq) + F–(aq)
• When acid is added, the equilibrium shifts to
the left.
Buffered Solutions (cont.)
• Additional H+ ions react with F– ions to form
undissociated HF molecules but the pH
changes little.
• The amount of acid or base that a buffer
solution can absorb without a significant
change in pH is called the buffer capacity.
Buffered Solutions (cont.)
• A buffer is most effective when the
concentrations of the conjugate acid-base
pair are equal or nearly equal.
Section 18.4 Assessment
In a neutralization reaction, an acid and
base react to form:
A. salt and oxygen gas
B. salt and ammonia
D
A
0%
C
D. precipitate and water
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. salt and water
Section 18.4 Assessment
Solutions that resist changes in pH are
called ____.
A. titrants
B. salts
D
A
0%
C
D. buffers
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. conjugate pairs
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Section 18.1 Introduction to Acids
and Bases
Key Concepts
• The concentrations of hydrogen ions and hydroxide
ions determine whether an aqueous solution is
acidic, basic, or neutral.
• An Arrhenius acid must contain an ionizable hydrogen
atom. An Arrhennius base must contain an ionizable
hydroxide group.
• A Brønsted-Lowry acid is a hydrogen ion donor. A
Brønsted-Lowry base is a hydrogen ion acceptor.
• A Lewis acid accepts an electron pair. A Lewis base
donates an electron pair.
Section 18.2 Strengths of Acids
and Bases
Key Concepts
• Strong acids and strong bases are completely ionized in
a dilute aqueous solution. Weak acids and weak bases
are partially ionized in a dilute aqueous solution.
• For weak acids and weak bases, the value of the acid
or base ionization constant is a measure of the strength
of the acid or base.
Section 18.3 Hydrogen Ions and pH
Key Concepts
• The ion product constant for water, Kw, equals the
product of the H+ ion concentration and the OH– ion
concentration.
Kw = [H+][OH–]
• The pH of a solution is the negative log of the hydrogen
ion concentration. The pOH is the negative log of the
hydroxide ion concentration. pH plus pOH equals 14.
pH = –log [H+]
pOH = –log [OH–]
pH + pOH = 14.00
• A neutral solution has a pH of 7.0 and a pOH of 7.0
because the concentrations of hydrogen ions and
hydroxide ions are equal.
Section 18.4 Neutralization
Key Concepts
• In a neutralization reaction, an acid and a base react
to form a salt and water.
• The net ionic equation for the neutralization of a strong
acid by a strong base is H+(aq) + OH–(aq) → H2O(l).
• Titration is the process in which an acid-base
neutralization reaction is used to determine the
concentration of a solution.
• Buffered solutions contain mixtures of molecules and
ions that resist changes in pH.
Which type of acid accepts an electron
pair?
A. Arrhenius
B. Brønsted-Lowry
D
A
0%
C
D. Dalton
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. Lewis
What is the conjugate of a weak acid?
A. strong acid
B. strong base
C. weak acid
D
C
A
0%
B
D. weak base
A. A
B. B
C. C
0%
0%
0%
D. D
In a solution with a pH of 4.0, which of the
following is true?
A. [H+] > [OH–]
B. [H+] < [OH–]
D
A
0%
C
D. none of the above
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. [H+] = [OH–]
Salt hydrolysis is what type of reaction?
A. single-replacement
B. double replacement
C. synthesis
D
C
A
0%
B
D. decomposition
A. A
B. B
C. C
0%
0%
0%
D. D
The strength of a weak acid is measured
by:
A. ion product constant
B. base ionization constant
D
A
0%
C
D. acid ionization constant
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. pOH
What is the pOH of 0.250M HBr, a strong
acid?
A. 0.60
B. 0.80
D
A
0%
C
D. 13.2
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. 13.4
Which of the following is a diprotic acid?
A. CH2COOH
B. HCl
D
A
0%
C
D. H2SO4
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. NH3+
What is the conjugate acid in the following
equation?
HCl(s) + H2O → H3O+(aq) + Cl–(aq)
A. Cl–
A
0%
D
D. HCl
C
C. H2O
A. A
B. B
C. C
0%
0%
0%
D. D
B
B. H3O+
An increase in concentration of hydrogen
ions when hydroxide ions are added to an
aqueous solution is an example of:
A. an Arrhenius acid
A
0%
D
D. Le Châtelier’s Principle
C
C. the equivalence point
A. A
B. B
C. C
0%
0%
0%
D. D
B
B. a Brønsted-Lowry acid
Which of the following is a polar
molecule?
A. CH4
B. PCl5
D
A
0%
C
D. H2
A. A
B. B
C. C
0%
0%
0%
D. D
B
C. H2O
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Table 18.2
Three Models for Acids and Bases
Table 18.3
Ionization Equations
Figure 18.20 A Neutralization Reaction
Figure 18.21 Titration
Figure 18.22 Neutralization Reactions
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