be able to explain the difference in base strength between ammonia,
primary aliphatic and primary aromatic amines in terms of the availability
of a lone pair on the N atom
According to Bronsted Lowry theory of acids and bases, bases are defined as
species that react with hydrogen ions in solution making a conjugate acid.
Hence Bronsted Lowry bases have a lone pair of electrons which can coordinate
to a hydrogen ion. Ammonia is a Bronsted Lowry base as are the organic derivatives
of ammonia, the amines.
The strength of a base is characterised by the availability of the lone pair
of electrons. When the electron density of the lone pair is high then the base
is able to coordinate easily to a hydrogen ion and it is a stronger base.
Amines are classified as primary, secondary or tertiary, depending on the number
of alkyl groups attached to the nitrogen atom. A primary amine (1º) has
only one alkyl group attached to the nitrogen atom, a secondary amine (2º)
has two, and a tertiary amine (3º) three.
Alkyl groups are electron inducing (they seem to act like little electron 'pumps',
pushing electrons away from themselves increasing the electron density on the
lone pair of the nitrogen. Hence the basicity of amines increases from 1º
to 2º to 3º.
Aromatic amines have the amine group directly attached to a benzene ring. Even
though the geometry is not perfect, the lone pair of electrons on the nitrogen
of the amine group is able to delocalise into the pi system of the sp2 carbon
ring. When a hydrogen ion coordinates to the lone pair on the nitrogen it reduces
the degree of delocalisation possible and hence destabilises the system.
This makes the basicity of aromatic amines lower than that of aliphatic (non-aromatic)
amines. For this reason phenylamine is a weaker base than methylamine.