Acids and Bases:
Introduction
Section 19.1
Objectives
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Identify the physical and chemical
properties of acids and bases
Classify solutions as acidic, basic, or neutral
Compare the Arrhenius and BronstedLowry models of acids and bases
Key Terms
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Acidic solutions
Basic solutions
Arrhenius model
Bronsted-Lowry model
Conjugate acid
Conjugate base
Conjugate acid-base pair
amphoteric
Properties of Acids and Bases
Acids
Bases
Tastes sour
Tastes bitter
Turns litmus red
Turns litmus blue
React with base to form
salt & water
Electrolyte
React with acid to form
salt & water
Electrolyte
React with metal to form Feels slippery
hydrogen gas
Self-Ionization
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Water molecules can react to form hydronium
(H3O+)and hydroxide ions(OH-)
H2O + H2O ↔ H3O+ + OH-

H2O ↔ H+ + OH-
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The relative amounts of hydronium ions (also
referred to as hydrogen ions, H+) and hydroxide
ions determine whether a solution is acid, base, or
neutral.
Determining acid, base,
neutral
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If hydrogen ions (H+)and hydroxide ions
(OH-) are equal in concentration, the
solution is neutral
If there are more hydrogen ions (H+) than
hydroxide ions (OH-), the solution is acidic
If there are more hydroxide ions (OH-), than
hydrogen (H+) ions the solution is basic
Arrhenius model of acid
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Arrhenius acid = a substance that contains
hydrogen and ionizes to produce hydrogen
ions in water.
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Example:
• HCl (aq)  H+ (aq) + Cl- (aq)
• HCl + H2O  H3O+ + Cl-
Arrhenius model of base
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Arrhenius base = a substance that contains
a hydroxide group and dissociates to
produce a hydroxide ions in aqueous
solution.
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Example:
• NaOH  Na+ + OH-
Arrhenius
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The Arrhenius acid and base model explains
many acids and bases.
However, there are acids and bases that
have the appropriate properties, but do not
fit the Arrhenius model.
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ENTER Brønsted-Lowry . . .
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Brønsted-Lowry Acid
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a hydrogen-ion donor
Example:
HCl + H2O → H3O+ + Cl –
HCl is an acid because it donates a
hydrogen ion to the water.
Of course, it was an acid according to
Arrhenius too. So what is different?
Brønsted-Lowry Acid
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Example:
NH3 + H2O ↔ NH4+ + OH-
Water is the acid here, because it donates a
hydrogen to the ammonia.
Brønsted-Lowry Base
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a hydrogen ion acceptor
NH3 + H2O ↔ NH4+ + OH-
Ammonia is the base here, because it accepts a
hydrogen from the water.
Brønsted-Lowry Base
HCl + H2O  H3O+ + Cl –
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Water is the base here because it accepts the
hydrogen ion from HCl.
Amphoteric
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a substance that can act as a BrønstedLowry acid OR base.
Example: water as an acid
NH3 + H2O ↔ NH4+ + OHExample: water as a base
HCl + H2O  H3O+ + Cl –
Conjugate acid-base pairs
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Conjugate acid = species produced when a
bases accepts a hydrogen ion from an acid
Conjugate base = species that results from
an acid donating a hydrogen ion to a base
NH3 + H2O ↔ NH4+ + OHbase acid
conj a conj b
More on page 599
Monoprotic and Polyprotic
Acids
(all classes)
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Monoprotic acid = acid that produces one
hydrogen ion (HCl).
Polyprotic acid = acid that produces more
than one hydrogen ion.
Diprotic acid = acid that produces 2
hydrogen ions (H2SO4)
Triprotic acid = acid that produces 3
hydrogen ions (H3PO4)
Strengths of Acids
and Bases
Section 19.2
Objectives
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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 and the
strength of a weak base with the strength of
its conjugate acid
Explain the relationship between the
strengths of acids and bases and the values
of their ionization constants
Key Terms
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Strong acid
Weak acid
Acid ionization constant
Strong base
Weak base
Base ionization constant
Strengths of Acids
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Strong acid = acid that ionizes completely.
Common examples:
HCl, HNO3, H2SO4
Others are listed on table 19-1
HCl  H + + Cl HCl + H2O  H3O+ + Cl –
Weak Acid
Weak acid = acid that does not ionizes
completely.
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Examples: acids other than those
memorized as strong.
HCN ↔ H+ + CN HCN + H2O ↔ H3O+ + CN -
Strengths of Bases
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Strong base = base that dissociates
completely to the metal ion and hydroxide
ion.
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Examples: Group I & II hydroxides
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NaOH(s)  Na+ + OH-
Weak Bases
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weak base = base that does NOT dissociates
completely to the metal ion and hydroxide
ion.
Examples: any base other than group I & II
hydroxides
NH3 + H2O ↔ NH4+ + OH-
Weak vs Strong Electrolyte
Concentration
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Is concentrated the same as strong?
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Is dilute the same as weak?
NO
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Concentrated is NOT the same as strong
when you are discussing acids and bases.
Dilute is NOT the same as weak when you
are discussing acids and bases.
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As you just learned, strong and weak have
to do with ionization or dissociation being
complete or incomplete.
In an earlier chapter, you learned that
concentrated and dilute have to do with the
amount of water added, or molarity of a
solution.
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Hydrochloric acid is ALWAYS a strong acid
because it ionizes completely.
However, it may be concentrated because
little to no water is added.
Or- it may be dilute if a lot of water is
added.
HCl can be a concentrated strong acid or a
dilute strong acid.
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Ammonia is ALWAYS a weak base because
it doesn’t dissociate completely.
However, it may be concentrated because
little to no water is added.
Or- it may be dilute if a lot of water is
added.
Ammonia can be a concentrated weak base
or a dilute weak base.
Acid ionization constant, Ka
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Just like Keq
HCN (aq) + H2O (l) ↔ H3O+ (aq) + CN- (aq)
Ka = [H3O] [CN]
[HCN]
Mathematically, a large Ka means larger
numerator more ions;
so the larger the Ka , the more the acid ionizes and
the stronger it is.
Acid ionization constant, Ka
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Table 19-2 page 605
Polyprotic acids have a Ka for each H
Base ionization constant, Kb
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Just like Ka
CH3NH3 (aq) + H2O (l) ↔ CH3NH4 +(aq) + OH- (aq)
Kb = [CH3NH4+][OH-]
[CH3NH3]
Mathematically, a large Kb means larger
numerator more ions;
so the larger the Kb , the more the base ionizes
and the stronger it is.