•According to the Arrhenius concept, a base is a
substance that produce OH- ions in aqueous solution.
•According to the Brønsted-Lowry model, a base is a
proton acceptor.
•A strong base dissociates completely when dissolved
in aqueous solution:
NaOH(s) → Na+(aq) + OH-(aq)
•All the hydroxides of the Group I elements (LiOH,
NaOH, KOH, RbOH, and CsOH) are strong bases.
•The alkaline earth (Group 2) hydroxides – Ca(OH)2,
Ba(OH)2, and Sr(OH)2 – are strong bases.
•Many types of proton acceptors (bases) do not
contain the hydroxide ion.
•However, when dissolved in water, these substances
increase the concentration of hydroxide ions because
of their reaction with water; thus, they yield a basic
solution.
•Note below the ammonia molecule accepts a proton
and thus functions as a base.
•The general reaction between a base B and water is
given by
B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq)
Base
Acid
Conjugate
acid
Conjugate
base
•The equilibrium constant for this general reaction is
[BH ][OH ]
Kb 
[B]
where Kb always refers to the reaction of a base with
water to form the conjugate acid and the hydroxide
ion.
•Bases of the type represented by B are called weak
bases.
•pH calculations for solutions of weak bases
are very similar to those for weak acids.
•Some important acids, such as sulfuric acid (H2SO4)
and phosphoric acid (H3PO4), can furnish more than
one proton and are called polyprotic acids.
•A polyprotic acid always dissociates in a stepwise
manner, one proton at a time.
•For example, the diprotic acid carbonic acid (H2CO3),
dissociates in the following steps:
H2CO3(aq) ⇌ H+(aq) + HCO3-(aq) Ka1 = 4.3 x 10-7
HCO3(aq) ⇌ H+(aq) + CO32-(aq)
Ka2 = 5.6 x 10-11
•Typically, successive Ka values are so much smaller than the first
value that only the first dissociation step makes a significant
contribution to the equilibrium concentration and the calculation of
pH for a solution of a typical weak polyprotic acid is identical to
that for a solution of a weak monoprotic acid.
•Sulfuric acid is unique in being a strong acid in its first
dissociation step and a weak acid in its second step. For
relatively concentrated solutions of sulfuric acid (1.0 M or higher),
the large concentration of H+ from the first dissociation step
represses the second step, which can be neglected as a
contributor of H+ ions. For dilute solutions of sulfuric acid, the
second step does make a significant contribution, and the
quadratic equation must be used.
•Salt is simply another name for ionic compound.
•When a salt dissolves in water, we assume it breaks
up into ions, which are independent units.
•Under certain conditions, these ions can behave as
acids or bases.
•Salts that consist of the cations of strong bases and
the anions of strong acids have no effect on [H+] when
dissolved in water.
•Aqueous solutions of salts such as KCl, NaCl,
NaNO3, and KNO3 are neutral (pH = 7).
•For example,
KCl(s)
→
K+(aq)
+
Cl-(aq)
cation of strong
base, KOH
anion of strong
acid, HCl
•In an aqueous solution of sodium acetate (NaC2H3O2),
the major species are
Na+, C2H3O2-, and H2O
•What are the acid-base properties of each component?
•The Na+ ion has neither acid nor base properties.
•The C2H3O2- ion is the conjugate base of acetic acid, a
weak acid. This means that C2H3O2- has a significant
affinity for a proton and is a base.
•Water is a weakly amphoteric substance.
•The pH of this solution will be determined by the
C2H3O2- ion, since it is a base.
•In this solution, water is the only source of protons to
react with the base, C2H3O2-, and the reaction is:
C2H3O2-(aq) + H2O(l) ⇌ HC2H3O2(aq) + OH-(aq)
•Since OH- is produced, Kb is defined as the
equilibrium constant; therefore,
[HC2H3O2 ][OH ]
Kb 

[C2H3O2 ]
•How can we obtain the Kb value for the acetate ion?
•The value of Ka for acetic acid is known (1.8 x 10-5)
and can be used along with Kw to obtain Kb for the
acetate ion.
•For any weak acid and its conjugate base,
Ka x Kb = Kw
•Thus, when either Ka or Kb is known, the other can be
calculated.
•For the acetate ion,
Kw
1.0 x 10 14
10
Kb 


5
.
6
x
10
K a ( for acetic acid ) 1.8 x 10 5
•Some salts produce acidic solutions when dissolved in
water.
•For example, when solid NH4Cl is dissolved in water,
NH4+ and Cl- ions are present, with NH4+ behaving as a
weak acid:
NH4+(aq) ⇌ NH3(aq) + H+(aq)
conjugate
acid
weak
base
•The Cl- ion does not affect the pH of the solution.
•In general, salts in which the anion is not a base and
the cation is the conjugate acid of a weak base produce
acidic solutions.
•A second type of salt that produces an acidic solution
is one that contains a highly charged metal ion.
•For example, when solid aluminum chloride (AlCl3) is
dissolved in water, the resulting solution is highly
acidic.
•Although the Al3+ ion is not itself a Brønsted-Lowry
acid, the hydrated ion Al(H2O)63+ formed in water is a
weak acid:
Al(H2O)63+(aq) ⇌ Al(OH)(H2O)52+(aq) + H+(aq)
•Typically, the higher the charge on the metal ion, the
stronger the acidity of the hydrated ion.
•For many salts, such as ammonium acetate
(NH4C2H3O2), both ions can affect the pH of the
aqueous solution.
•Compare Ka and Kb values for the acidic and basic
ions and predict based on table below.