Chemistry II Notes
Chapter 16: Acids & Bases
Acid-Base Equilibria
16.1. Acids and Bases: A Brief Review
Acid and base behavior in an aqueous medium is briefly reviewed.
Arrhenius defined an acid as a substance that produces H+ ions in water; he
defined a base as a substance that produces OH– ions in water. HCl—one of the
strong acids—is an Arrhenius acid. Sodium hydroxide—one of the strong bases—
is an Arrhenius base.
16.2 Brønsted-Lowry Acids and Bases
The Arrhenius definitions of acids and bases are limited to aqueous reactions. A
broader definition is the Brønsted-Lowry approach. According to Brønsted-Lowry,
an acid is a substance that donates an H+ ion to another substance; a base is a
substance that accepts an H+ ion.
Chemists use the term proton to refer to the aqueous hydrogen ion, H+. They also
refer to the hydronium ion, H3O+ in the context of acids. All four of these terms,
H+, hydrogen ion, H3O+, and hydronium ion, are used interchangeably. In water a
proton interacts with the lone pairs on oxygen atoms of water molecules and
becomes hydrated. So it is slightly more realistic to represent an aqueous proton
as H3O+, which depicts the proton attached to a water molecule.
In order for one substance to behave as an acid, another substance must behave as
a base (and vice versa). A Brønsted-Lowry acid must have a proton to donate, and
a Brønsted-Lowry base must have a lone pair of electrons in order to accept the
proton. Some substances are capable of acting as an acid in one reaction and as a
base in another. Such substances are called amphoteric.
When an acid molecule ionizes in water, it donates its proton to a water molecule,
producing a hydronium ion and an anion. The anion produced by the ionization is
the conjugate base of that acid. (HX is used to denote a generic acid.)
HX (aq) + H2O (l)  H3O+ (aq) + X- (aq)
HX (the acid) and X– (the conjugate base) differ only by a proton. They are a
conjugate acid-base pair. Every acid has a conjugate base. By the same token,
every base has a conjugate acid.
16.3 The Autoionization of Water
Pure water has a very small tendency to ionize, acting both as an acid (donating a
proton) and as a base (accepting a proton).
At 25°C, the Kc for this process is 1.0 10–14, which means that only about one
molecule per billion undergoes this autoionization. The equilibrium expression for
the autoionization of water is
(Recall that a liquid does not appear in the equilibrium expression.)
Because the autoionization of water is a very important equilibrium, its equilibrium
constant is given a special subscript, w. For any aqueous solution at 25°C, the
product of hydronium and hydroxide ion concentrations is equal to Kw. In neutral
water, where the only source of either ion is the autoionization, the hydronium and
hydroxide ion concentrations are equal.
So the concentrations of both hydronium ion and hydroxide ion in neutral water is
1.0 10–7 M.
16.4 The pH Scale
16.5 Strong Acids and Bases
Strong Acids
Hydrochloric acid
Hydrobromic acid
Hydroiodic acid
Strong Bases
HCl
LiOH
Lithium hydroxide
HBr
NaOH
Sodium hydroxide
HI
KOH
Potassium
hydroxide
HNO3
Nitric acid
RbOH
Rubidium
hydroxide
HClO3
Chloric acid
CsOH
Cesium hydroxide
HClO4
Perchloric acid
Ba(OH)2
Barium hydroxide
H2SO4
Sulfuric acid
Sr(OH)2
Strontium
hydroxide
Strong acids are those that ionize completely in water. Strong acids are also strong
electrolytes.
Strong bases are ionic compounds that dissociate completely in water. They
include the hydroxides of group 1A metals and group 2A metals.
16.6 Weak Acids
A weak acid is one that ionizes partially in water to produce hydronium ion and a
conjugate base. Acetic acid is a weak acid.
The ionization equilibrium of a weak acid has an associated equilibrium constant
called the acid-dissociation constant. The equilibrium constant for a weak acid
ionization is subscripted with an a for acid. The equilibrium expression for the
above equation is
As with all equilibria, the larger the value of K, the further the equilibrium lies to
the right. This means that the larger the value of Ka, the stronger the acid. Table
16.2 lists some weak acids, their conjugate bases, and their acid-dissociation
constants.
16.7 Weak Bases
A weak base is one that ionizes partially in water to produce hydroxide ion and a
conjugate acid. Ammonia is an example of a group of weakly basic compounds
called amines. An amine contains a nitrogen atom with a lone pair of electrons.
The site of the lone pair is where a proton can be accepted. In general,
Being an equilibrium, the ionization of a weak base has an equilibrium constant
called the base-dissociation constant associated with it. The equilibrium constant
for a weak base ionization has the subscript b for base. The equilibrium expression
for the above equation is
As with all equilibria, the larger the value of K, the further the equilibrium lies to
the right. This means that the larger the value of Kb, the stronger the base. Table
16.4 lists some weak bases, their conjugate acids, and their base-dissociation
constants.
16.8 Relationship Between Ka and Kb
As we saw in Section 16.2, a weak acid has a strong conjugate base. It is important
to distinguish between the meaning of the word strong in the context of acids and
bases and its meaning in the context of conjugate acids and bases. A strong acid is
one that ionizes completely in water. A strong conjugate acid is one that is
sufficiently acidic compared with water to protonate the water to some extent. Note
again the difference in equations used to represent each. The ionization of HCl is
written with an arrow in one direction. The ionization of the ammonium ion (the
conjugate acid of the weak base ammonia) shows a double arrow, denoting an
equilibrium.
16.9 Acid-Base Properties of Salt Solutions
The pH of a salt solution can be predicted qualitatively by considering the ionic
constituents of the salt.
1. Salts derived from a strong base and a strong acid: Examples are NaCl and
Ca(NO3)2, which are derived from NaOH and HCl and from Ca(OH)2 and
HNO3, respectively. Neither cation nor anion hydrolyzes. The solution has a
pH of 7.
2. Salts derived from a strong base and a weak acid: In this case the anion is a
relatively strong conjugate base. Examples are NaClO and Ba(C2H3O2)2.
The anion hydrolyzes to produce OH–(aq) ions. The cation does not
hydrolyze. The solution has a pH above 7.
3. Salts derived from a weak base and strong acid: In this case the cation is a
relatively strong conjugate acid. Examples are NH4Cl and Al(NO3)3. The
cation hydrolyzes to produce H+(aq) ions. The anion does not hydrolyze.
The solution has a pH below 7.
4. Salts derived from a weak base and a weak acid: Examples are NH4C2H3O2,
NH4CN, and FeCO3. Both cation and anion hydrolyze. The pH of the
solution depends on the extent to which each ion hydrolyzes. The pH of a
solution of NH4CN is greater than 7 because CN– (Kb = 2.0 10–5) is more
basic than NH4+ (Ka = 5.6 10–10) is acidic. Consequently, CN– hydrolyzes
to a greater extent than NH4+ does.
16.10 Acid-Base Behavior and Chemical Structure
As we move from left to right across a row in the periodic table, there is less
change in bond strength. In this case, what determines acid strength is the polarity
of the H- X bond. The electronegativity of elements increases from left to right
across a period in the periodic table. As the electronegativity of X increases, the
polarity of the H- X bond increases, increasing acidity.
The strength of an oxyacid depends on the electronegativity of the central
nonmetal to which the OH groups are bound and on the number of oxygen atoms
bound to the central nonmetal atom. For a series of oxyacids with the same number
of oxygen atoms, the acidity increases with the electronegativity of the nonmetal.
For a series of acids with the same central nonmetal atom, the acidity increases
with the number of oxygen atoms bound to the central atom. (This also relates
increasing acidity to increasing oxidation number on the central atom.)
Carboxylic acids are all weak acids. Their acidity stems partly from the fact that
the second oxygen withdraws electron density from the O–H bond, making it more
polar than it would otherwise be. The acidity is increased by the presence of
additional electron-withdrawing (electronegative) atoms, such as halogens. In
addition, the anion left behind after donation of a proton (the conjugate base) is
fairly stable. The negative charge on the anion can be accommodated on either of
the two oxygen atoms, stabilizing the anion via resonance.
16.11 Lewis Acids and Bases
The Brønsted-Lowry acid-base theory broadens the definitions of acids and bases
to include reactions that do not occur in aqueous media. Lewis acid-base theory
further broadens the definition to include reactions other than proton-transfer
reactions.
A Lewis acid is defined as an electron-pair acceptor, and a Lewis base as an
electron-pair donor. In the examples we have seen of Arrhenius and BrønstedLowry acid-base behavior, the bases are all acting as Lewis bases. In the ionization
of an acid in water, the water molecule donates electrons (a lone pair on the
oxygen atom) to the hydrogen atom of the acid. As a bond forms between the
oxygen atom on water and the acidic proton, the bond between the acidic proton
and its original molecule is broken.
The Lewis definition of an acid does not require that an acid have a proton to
donate, only that it be able to accept a pair of electrons from a Lewis base.