The use of acid and base as classifications of chemical
substances is an ancient one and predates even alchemy. Acids
were originally substances that tasted sour and reacted with
bases; bases were alkaline salts such as CaO (lime), NaOH, and
KOH, while amphoteric substances were substances which could
react with either acids or bases. In modern chemistry, these
concepts have taken on considerably more precise meaning.
The chemistry of acids and bases is the chemistry of
reactions in which protons are transferred.
Arrhenius Theory of Acid-Base Behavior
The first quantitative approach to acid-base equilibria was
developed around 1884 by the Swedish chemist Svante Arrhenius.
According to his Arrhenius theory, an acid is any compound or
ion which yields hydrogen ions in water solution and a base is
any compound or ion which yields hydroxyl ions in water
The approach taken by Arrhenius with acids and bases was in
accordance with his general approach to substances in aqueous
solution. Materials which dissolved in whole or in part as ions
were, and still are, called electrolytes. A
strong electrolyte, like sodium chloride, was a
substance which dissociated completely into ions while a
weak electrolyte remained partially in molecular form.
Thus a strong acid is a strong electrolyte
which ionized completely to give hydrogen ions in aqueous
solution. All of the common strong acids - sulfuric acid, H2SO4;
nitric acid, HNO3; hydrochloric acid, HCl; and
perchloric acid, HClO4 - easily fit in this category.
Likewise, a strong base is a strong electrolyte
which ionized completely to give hydroxyl ions in aqueous
solution. The common strong bases, NaOH and KOH, fit in this
category. Acid-base reactions are then simply the reaction of
the hydrogen ion with the hydroxyl ion as a
neutralization reaction, H+(aq) + OH-(aq)
--> H2O. The remaining ions simply continue to exist
as a solution of a salt. The reaction including these "bystander
ions" could be written as
H+(aq) + Cl-(aq) + Na+(aq) +
OH-(aq) --> Na+(aq) + Cl-(aq) +
The Arrhenius concept was not intended to, and does not, deal
with acid-base chemistry in solvents other than water. Even in
water, however, it has deficiencies. Ammonia, NH3, clearly
behaves in a qualitative manner as a base and reacts to
neutralize acids. However, ammonia does not contain a hydroxide
ion. This is of itself no obstacle because ammonia might well
produce hydroxyl ion by reaction with water:
NH3(aq) + H2O --> NH3.HOH
--> NH4OH --> NH4+(aq) + OH-(aq)
However, no experimental evidence for the existence of NH4OH
has yet been foumd. For these two reasons, the Arrhenius concept
is no longer considered an adequate treatment of acids and bases
even in aqueous solution.
Bronsted-Lowry Theory of Acid-Base Behavior
The Arrhenius theory was the only theory used to explain the
behaviour of acids and bases for about forty years. In 1923, J.
N. Bronsted in Denmark and T. M. Lowry in England independently,
and almost simultaneously, proposed the modern "protonic" or
Bronsted-Lowry theory of acid-base behaviour. According to the
Bronsted-Lowry concept, an acid is any compound or ion
which can give up a proton, while a base is any compound or ion
which can accept a proton.
Any species which can either accept or give up a proton is
said to be amphiprotic. Thus the water molecule
is amphiprotic, since it can give up a proton, H2O
--> H+ + OH-, to form the hydroxyl ion OH-.
Alternatively, water can accept a proton to form the
hydronium ion H3O+, according to
the equation H+ + H2O --> H3O+.
The above two equations can be combined to give the dissociation
equation for water: 2H2O --> H3O+
The Bronsted-Lowry concept is an extension of the Arrhenius
concept in that all Arrhenius bases, being sources of hydroxide,
can accept protons. Ammonia and amines will also accept protons
to form the corresponding ammonium ions, so the existence of
NH4OH is no longer necessary to explain the action of ammonia as
a base. The Bronsted-Lowry concept also is useful in protonic
solvents other than water, such as liquid ammonia or glacial
acetic acid, where the Arrhenius concept is not useful. We will,
however, generally confine our discussion to aqueous solutions
because they are so much more important.
Lewis Explanation of Acid-Base Behavior
The basic principles of the Lewis theory of acid-base behaviour
were also set down in 1923, by the American physical chemist G.
N. Lewis. The Lewis definitions of acids and bases are even more
inclusive than the Bronsted definitions. The Lewis definitions
are that an acid is an electron-pair acceptor and a base
is an electron-pair donor.
Since a base like ammonia, H3N:, has a lone pair
of electrons, it can be considered to "donate" them to a proton
in forming the conjugate acid NH4+. The
Lewis definitions are used to explain the effect of compounds
such as AlCl3, which acts as an acid in nonaqueous
organic solvents, on organic reactions. In protonic solvents,
however, the Lewis definitions are far less useful than are the