Chapter 19
 All
aqueous solutions contain hydrogen
ions (H+) and hydroxide (OH-) ions.
 An acidic solution contains more H+ ions
than OH-.
 A basic solution contains more OH- ions
than H+ ions.
 When a solution has the same
concentration of H+ and OH- it is said to
be neutral.
A
hydronium ion (H3O+) is a hydrated
hydrogen ion.
 H+ and H3O+ mean the same thing and
can be used interchangeably.
 Arrhenius
model
• An acid is a substance that contains hydrogen
and ionizes in aqueous solution to produce
hydrogen ions.
• A base is a substance that contains a hydroxide
group and dissociates in aqueous solution to
produce hydroxide ions.
 Brønsted-Lowry
model
• An acid is a hydrogen-ion donor.
• A base is a hydrogen-ion acceptor.
 When
a Brønsted-Lowry acid donates a
hydrogen ion, a conjugate base is
formed.
 When a Brønsted-Lowry base accepts a
hydrogen ion, a conjugate acid is
formed.
 Identify
the conjugate acid-base pairs in
this reaction.
HClO2 (aq) + H2O (l)  H3O+ (aq) + ClO2- (aq)
1.
Indentify the conjugate acid-base pairs in
the following reactions.
a. H2SO3 (aq) + H2O (l)  HSO3- (aq) + H3O+ (aq)
b. HPO42- (aq) + H2O (l)  H2PO4- (aq) + OH- (aq)
c. HSeO3- (aq) + H2O (l)  H3O+ (aq) + SeO32- (aq)
 An
acid that can donate only one
hydrogen ion is called a monoprotic acid.
 For example, hydrochloric acid (HCl) and
formic acid (HCOOH).
 Note: Only those hydrogens that are
bonded to electronegative elements are
ionizable.
 Some
acids can donate more than one
hydrogen ion.
 For example, sulfuric acid (H2SO4)
contains 2 ionizable hydrogen atoms, and
is called a diprotic acid.
 Boric acid (H3BO3) contains 3 ionizable
hydrogen atoms, and is called a triprotic
acid.
 An acid with two or more ionizable
hydrogens is called a polyprotic acid.
 The
three ionizations of boric acid are as
follows.
H3BO3 (aq) + H2O (l)  H3O+ (aq) + H2BO3- (aq)
H2BO3- (aq) + H2O (l)  H3O+ (aq) + HBO32(aq)
HBO3 2-(aq) + H2O (l)  H3O+ (aq) + BO3 3- (aq)
2.
Write the steps in the complete
ionization of the following polyprotic
acids.
a. Carbonic acid (H2CO3)
b. Chromic acid (H2CrO4)
 An
acid that ionizes completely in dilute
aqueous solution is called a strong acid.
 There are six strong acids that you need
to learn, and they are listed on page 603
in your textbook.
 A weak acid is one that ionizes only
partially in dilute aqueous solutions.
 For a weak acid, a state of equilibrium is
reached in which the forward and reverse
reactions occur at equal rates.
 Considering
the reaction of the weak
acid formic acid
HCOOH (aq) + H2O (l)  H3O+ (aq) + HCOO- (aq)
 The
equilibrium constant expression of
the ionization of formic acid in water is as
follows:
Ka= [H3O+][HCOO-]
[HCOOH]
 Ka is the acid ionization constant.
 Ka
is the value of the equilibrium constant
for the ionization of a weak acid.
 Ka is a measure of the extent of ionization
of the acid.
 Weak acids have the smallest Ka values.
 Polyprotic
acids have a Ka value for each
ionization, and the Ka values decrease for
each successive ionization.
3.
Write ionization equations and acid
ionization constant expressions for the
following acids.
a. Hydrofluoric acid (HF)
b. Hypobromous acid (HBrO)
4.
Write the ionization equation and the
acid ionization constant expression for
the second ionzation of sulfurous acid
(H2SO3) in water.
 Metallic
hydroxides are strong bases
which dissociate entirely into metal ions
and hydroxide ions in aqueous solution.
 Group 1A and 2A hydroxides are strong
bases.
 A weak base is a base that ionizes only
partially in dilute solution to form an
equilibrium mixture.
 Just
like there is an acid ionization
constant there is a base ionization
constant.
 The Kb value is the value of the
equilibrium constant.
 Kb is smallest for the weakest bases.
5.
Write ionization equations and base
ionization constant expressions for the
following bases.
a. Butylamine (C4H9NH2)
b. Phosphate ion (PO43-)
c. Hydrogen carbonate ion (HCO3-)
 Pure water self-ionizes slightly to form
H3O+ and OH- ions, as shown
H2O (l) + H2O (l)  H3O+ (aq) + OH- (aq)
 It can be simplified by removing one
water molecule from each side
H2O (l)  H+ (aq) + OH- (aq)
A
special equilibrium expression for the
self-ionization of water is defined as
follows:
Kw = [H+][OH-]
 Kw is called the ion product constant for
water.
 It is the value of the equilibrium constant
expression of water.
 In
pure water at 298 K, the concentration
of H+ ions and OH- ions both equal 1.0 x
10-7 M, so the value of Kw = 1.0 x 10-14.
 At
298 K, the OH- ion concentration of an
aqueous solution is 1.0 x 10-11 M. Find the
H+ ion concentration in the solution and
determine whether the solution is acidic,
basic or neutral.
6.
Given the concentration of either
hydrogen ion or hydroxide ion,
calculate the concentration of the other
ion at 298 K and state whether the
solution is acidic, basic, or neutral.
a. [OH-] = 1.0 x 10-6 M
b. [H+] = 1.0 x 10-7 M
c. [H+] = 8.1 x 10-3 M
 Because
the concentrations of H+ ions
are often small, the pH scaled was
developed.
 The pH of a solution equals the negative
logarithm of the hydrogen ion
concentration.
pH = -log [H+]
 The pH scale has values from 0 to 14.
 Acids have pHs less than 7 and bases
have pHs greater than 7.
A
pH of 0 is the most acidic and a pH of
14 is the most basic.
 A pH of 7 is neutral.
 The pOH scale expresses the basicity of a
solution.
 pOH is the negative logarithm of the
hydroxide ion concentration.
pOH = -log [OH-]
 If
either the pH or pOH are known, the
other may be determined by using the
following relationship.
pH + pOH = 14.00
 The pH and pOH values can be
determined if eith the [H+] or [OH-] is
known.
 If
a certain carbonated soft drink has a
hydrogen ion concentration of 7.3 x 10-4
M, what are the pH and pOH of the soft
drink?
7.
Calculate the pH and pOH of aqueous
solutions having the following ion
concentrations.
a. [H+] = 1.0 x 10-14 M
b. [OH-] = 5.6 x 10-8 M
c. [H+] = 2.7 x 10-3 M
d. [OH-] = 0.061 M
 When
the pH of a solution is known, you
can determine the concentrations of H+
and OH-.
 What
are [H+] and [OH-] in an antacid
solution with a pH of 9.70?
8.
The pH or pOH is given for three
solutions. Calculate [H+] and [OH-] in
each solution.
a. pH = 2.80
b. pH = 13.19
c. pOH = 8.76
 Remember, that
strong acids and bases
dissociate completely in water.
 This means that for monoprotic acids the
concentration is equal to the
concentration of the hydrogen ion.
 In some acids and bases, there are more
than one hydrogen ions or hydoxide ions
in the compound.
9.
Calculate the pH of the following strong
acid or strong base solutions.
a. 0.015 M HCl
b. 0.65 M KOH
c. 2.5 x 10-4 M HNO3
d. 4.0 x 10-3 M Ca(OH)2
 If
you know the pH and the concentration
of a solution of a weak acid, you can
calculate Ka for the acid.
 The
pH of a 0.200 M solution of acetic
acid (CH3COOH) is 2.72. What is the Ka
for acetic acid?
10.
Calculate Ka for the following acids
using the information provided.
a. 0.100M solution of sulfurous acid (H2SO3),
pH = 1.48
b. 0.200M solution of benzoic acid (C6H5COOH),
pH = 2.45
 The
reaction of an acid and a base in an
aqueous solution is called a
neutralization reaction.
 The products of a neutralization reaction
are always a salt and water.
 A salt is an ionic compound composed of
a positive ion from a base and a negative
ion from an acid.
 Acid-base
neutralizations are used in the
procedure called titration, which is a
method for determining the
concentration of a solution by reacting it
with another solution of known
concentration.
 Neutralization reactions proceed until an
equivalence point is reached.
 The equivalence point is the point where
the moles of H+ ions and OH- ions are
equal.
 At
the equivalence point, a large change
in pH occurs that can be detected by a
pH meter or an acid-base indicator.
 An indicator is a dye whose color is
affected by pH changes.
 When a strong acid is titrated with a
strong base, the equiv. point is 7.
 When strong acid, weak base equiv. point
is less than 7; when weak acid and strong
base the point is greater than 7.
 In
a titration, 53.7 mL of 0.100 M HCl
solution is needed to neutralize 80.0 mL
of KOH solution. What is the molarity of
the KOH solution?
11.
A 45.0 ml sample of nitric acid solution
is neutralized by 119.4 mL of 0.200 M
NaOH solution. What is the molarity of
the nitric acid solution?
A
buffer is a solution that resists changes
in pH when moderate amounts of acid or
base is added.
 A buffer is a mixture of a weak acid and
its conjugate base, or a weak base and its
conjugate acid.