Chapter 17: Acids and Bases
• Acid-base reactions involve proton
(hydrogen ion, H+) transfer
• The generalization of the Arrhenius
definition of acids and bases is called the
Brønsted-Lowry definitions:
• An acid is a proton donor
• A base is a proton acceptor
• This allows for gas phase acid-base
reactions
The reaction of HCl and H2O. HCl is the acid because it donates
a proton. Water is the base because it accepts a proton.
• Species that differ by a proton, like H2O and
H3O+, are called conjugate acid-base pairs
(a) Formic acid transfers a proton to a water molecule.
HCHO2 is the acid and H2O is the base. (b) When a
hydronium ion transfers a proton to the CHO2- ion, H3O+ is
the acid and formate ion is the base.
• An amphoteric substances can act as either
an acid or base
• These are also called amphiprotic, and can
be either molecules or ions
– For example, the hydrogen carbonate ion:
As an acid :
HCO3 (aq )  OH  ( aq )  CO32 (aq )  H 2O
As a base :
HCO3 (aq )  H 3O  ( aq )  H 2CO(aq )  H 2O
• The strength of an acid is a measure of its
ability to transfer a proton
• Acids that react completely with water (like
HCl and HNO3) are classified as strong
• Acids that are less than completely ionized
are called weak acids
• Bases can be classified in a similar fashion:
– Strong bases, like the oxide ion, react
completely
– Weak bases, like NH3, undergo incomplete
reactions
• The strongest acid in water is the
hydronium ion
• If a more powerful proton donor is added to
water, it quantitatively reacts with water to
produce H3O+
• Similarly, the strongest base that can be
found in water is the hydroxide ion, because
more powerful proton acceptors react
quantitatively with water to produce OH-
• Acetic acid (HC2H3O2) is a weak acid
• It ionizes only slightly in water


HC 2 H 3O2 (aq)  H 2O 
H
O
(
aq
)

C
H
O

3
2 3 2 ( aq )
weaker acid
we aker base
stronger acid
stronger base
– The hydronium ion is a better proton donor
than acetic acid (it is a stronger acid)
– The acetate ion is a better proton acceptor than
water (it is a stronger base)
• The position of an acid-base equilibrium
favors the weaker acid and base
• This can be generalized:
– Stronger acids and bases tend to react with each
other to produce their weaker conjugates
– The stronger a Brønsted acid is, the weaker is
its conjugate base
– The weaker a Brønsted acid is, the stronger is
its conjugate base
• These ideas can be applied to the binary
acids (acids made from hydrogen and one
other element)
• The strengths of the binary acids increases
from left to right within the same period
• For example, HCl is stronger acid than H2S which is
a stronger acid than PH3
• The strengths of the binary acids increase
from top to bottom within a group
• For example, HI is a stronger acid than HBr which
is a stronger acid than HCl
• Trends are also present in the oxoacids
(acids of hydrogen, oxygen, and one other
element)
• When the central atom holds the same
number of oxygen atoms, the acid strength
increases from the bottom to top within a
group and from left to right within a period
• Acid strength: HClO4 > HBrO4 > HIO4
• Acid strength: HClO4 > H2SO4 > H3PO4
• For a given central atom, the acid strength
of an oxoacid increases with the number of
oxygens held by the central atom
• Acid strength: H2SO4 > H2SO3
• The strength of an acid can be analyzed in
terms the the basicity of the anion formed
during the ionization
• The basicity is the willingness of the anion
to accept a proton from the hydronium ion
• Consider H2SO4 and H3PO4:
O
O
||
||
HO  S  OH
HO  P  OH
O
sulfuric acid
OH
phosphoric acid
||
|
• The anions are:

O


||


 HO  S||  O 


O



O


||


 HO  P|  O 


OH




HSO4 (3 lone O) H 2 PO4 (2 lone O)
• In oxoanions, the lone oxygens carry most
of the negative charge, making the
hydrogen sulfate ion a weaker base than
the hydrogen phosphate ion
• In terms of the percentage of molecules that
are ionized, sulfuric acid is a stronger acid
than phosphoric acid
• There is a third definition for acid and bases
• It is a further generalization, or broadening,
of the species that can be classified as either
an acid or base
• The definitions are based on electron pairs
and are called Lewis acids and bases
• A Lewis acid is any ionic or molecular
species that can accept a pair of electrons in
the formation of a coordinate covalent bond
• A Lewis base is any ionic or molecular
species that can donate a pair of electrons in
the formation of a coordinate covalent bond
• Neutralization is the formation of a
coordinate covalent bond between the donor
(base) and acceptor (acid)
NH3 (a Lewis base) forms a coordinate covalent bond with
BF3 (a Lewis acid) during neutralization. NH3BF3 is called an
addition compound because it was made by joining two
smaller molecules.
Carbon dioxide (a Lewis acid) reacts with hydroxide ion (a
Lewis base) in solution to form the bicarbonate ion. The
electrons in the coordinate covalent bond come from the
oxygen atom in the hydroxide ion.
• Lewis acids:
– Molecules or ions with incomplete valence
shells (for example BF3 or H+)
– Molecules or ions with complete valence shells,
but with multiple bonds that can be shifted to
make room for more electrons (for example
CO2)
– Molecules or ions that have central atoms
capable of holding additional electrons (usually,
atoms of elements in Period 3 and below, for
example SO2)
• Lewis bases:
– Molecules or ions that have unshared pairs of
electrons and that have complete shells (for
example O2- or NH3)
• All Brønsted acids and bases are Lewis
acids and bases, just like all Arrhenius acids
and bases are Brønsted acids and bases
• Consider a proton transfer from the Lewis
perspective
• For example, the proton transfer between
the hydronium ion and ammonia:

H 2O  H  NH 3  H 2O  NH

4
• The elements most likely to form acids are
the nonmetals in the upper right-hand
corner of the periodic table
• The elements most likely to form basic
hydroxides are the IA and IIA metals along
the left of the periodic table
• The elements themselves can be classified
according to the ability of their oxides to
form acids or bases
• In general, most metal oxides react with
water to form bases, and nonmetal oxides
react with water to form acids
• In Section 5.5 metal oxides were called base
anhydrides and nonmetal oxides were called acid
anhydrides
• When cations dissolve in water, they form
species called hydrated ions
• Hydrated metal ions tend to be Brønsted
acids
• For the monohydrate of the metal ion Mn+
the equilibrium can be represented as

M ( H 2O)n  H 2 O
MOH ( n1)  H 3O
The metal ion
makes the
hydrogens on
the water
more acidic.
• The charge density of a cation is its charge
divided by its volume
• The higher the charge density, the better a
cation is at drawing electron density from a
O-H bond and the more acidic it is
• Within a given period, the cation size increases, and
the charge density decreases, from top to bottom
• As a result, the most acidic hydrated cations are
found at the top of a group
• As the cation charge increases, it becomes
more acidic
• When the charge (oxidation number) is
small, its oxide tends to be basic
• When the cation ion charge is +3, the oxide
begins to become acidic
• An amphoteric species is capable of acting
as both an acid and base
• Aluminum oxide is an example of an
amphoteric compound
Al2O3 as an acid :
Al2O3(s)  2OH -(aq)  2 AlO2  H 2O
Al2O3 as a base :
Al2O3(s)  6 H (aq)  2 Al 3 (aq )  3H 2O
• Many nonmetal oxides are acid anhydrides
– For example:
SO3 ( g )  H 2O  H 2 SO4 (aq)
sulfuric acid
CO2 ( g )  H 2O  H 2CO3 (aq)
carbonic acid
• Water undergoes self-ionization or
autoionization making it a weak electrolyte



H 2O  H 2O H 3O  OH

• This equilibrium is described by the ion
product of water




K w  [ H 3O ][OH ]  [ H ][OH ]
• At 25oC in pure water it has been found that
[ H  ]  [OH  ]  1.0 107 M
so that
14
K w  1.0 10
(at 25 C )
• In any aqueous solution, the product of [H+]
and [OH-] equals Kw
• This provides an alternate a way to define
the acidity or basicity of a solution
o
• Neutral solutions: [H3O+] = [OH-] or [H+] = [OH-]
• Acidic solutions: [H3O+] > [OH-] or [H+] > [OH-]
• Basic solutions: [H3O+] < [OH-] or [H+] < [OH-]
• To make the comparison of small values of
[H+] easier, the pH was defined:


pH   log[ H ] or [ H ]  10
• In terms of the pH:
• Neutral solutions: pH = 7.00
• Acidic solutions: pH < 7.00
• Basic solutions: pH > 7.00
-pH
The pH of some
common solutions.
[H+] decreases,
while [OH-]
increases, from top
to bottom.
• The pH of a solution can be measured with
a pH meter or estimated using a visual acidbase indicator (see Table 17.3, page 762)
• An acid-base indicator is a species that
changes color based on the pH
• Calculating the pH of a strong acid or base
is “easy” because they are 100% dissociated
in aqueous
• For example, the pH of 0.10 M HCl is 1.00 and the
pH of 0.10 M NaOH is 13.00
• In the last example it was assumed that the
total concentration of [H+] was due to the
strong acid (HCl) and [OH-] was due to the
strong base (NaOH)
• This assumption is valid because the
autoionization of water is suppressed in
strongly acidic or strongly basic solutions
– This assumption fails for very dilute solutions
of acids or bases (less than 10-6 M)