A more Satisfactory theory was proposed in 1923 by the Danish chemist Johannes Bronsted-Lowry
(1874-1936), a British chemist. Their theory states that an acid is a proton
(hydrogen ion, H ) donor and a base a proton acceptor. Although the acid
must still contain hydrogen, the Bronsted-Lowry theory does not require
an aqueous medium. For example, liquid ammonia, which acts as a base in
aqueous solution, can act as an acid in the absence of water by transferring
a proton to a base and forming the amide anion (negative ion) NH2-. The
Bronsted-Lowry definition of acids and bases also explains why a strong
acid displaces a weak acid from its compounds (and likewise for strong and
weak bases). Here acid-base -reactions are viewed as a competition for protons.
In terms of a general chemical equation, the reaction of Acid (1) with Base
(2) results in the transfer of a proton from Acid (1) to Base (2). In losing
the proton, Acid (1) becomes its conjugate base, Base (1). In gaining a
proton, Base (2) becomes its conjugate acid, Acid (2). The equilibrium represented
by the equation above may be displaced either to the left or to the right,
and the actual reaction will take place in the direction that produces the
weaker acid-base pair. For example, HCl is a strong acid in water because
it readily transfers a proton to water to form a hydronium ion. The equilibrium
lies mostly to the right because the conjugate base of HCl, Cl -, is a weak
base, and H3O+, the conjugate acid of H2O, is a weak acid. In contrast,
hydrogen fluoride, HF, is a weak acid in water because it does not readily
transfer a proton to water. This equilibrium lies mostly to the left because
H2O is a weaker base than F, and because HF is a weaker acid (in water)
than H3O+. The Bronsted-Lowry theory also explains why water can be amphoteric,
that is, why it can serve as either an acid or a base. Water serves as a
base in the presence of an acid that is stronger than water (such as HCl),
in other words, an acid that has a greater tendency to dissociate than does
water. Water can also serve as an acid in the presence of a base that is
stronger than water (such as ammonia).
How can we measure Acid or Base Strength?
The strength of an acid can be measured by the extent to which an acid
transfers a proton to water to produce the hydronium ion, H3O+. Conversely,
the strength of a base is indicated by the extent to which the base removes
a proton from water.
A convenient acid-base scale is calculated from the amount of H3O+ that
is formed in water solutions of acids or of OH- formed in water solutions
of bases. The former is known as the pH scale and the latter as the pOH
scale ( see pH ). The value for pH is equal to the negative logarithm
(q.v.) of the hydronium ion concentration-and for pOH, of the hydroxyl
ion concentration-in an aqueous solution. Pure water has a pH of 7.0.
When an acid is added, the hydronium ion con centration [H3O+ ] becomes
larger than that in pure water, and the pH becomes less than 7.0, depending
on the strength of the acid. The pOH of pure water is also 7.0, and in
the presence of a base, the pOH drops to values lower than 7.0. The American
chemist Gilbert N. Lewis has offered another theory of acids and bases
that has the further advantage of not requiring the acid to contain hydrogen.
This theory states that acids are electron-pair acceptors and bases are
electron-pair donors. This theory also has the advantages that it works
when solvents other than water are involved and it does not require the
formation of a salt or of acid-base conjugate pairs. Thus, ammonia is
viewed as a base because it can donate an electron pair to the acid boron
trifluoride, for example. to form an acid-base association pair.
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