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Explain the basicity of aromatic amine and ammonia\[?\]

Last updated date: 17th Jul 2024
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Hint: The less electronegative the element, the less stable the lone pair will be and therefore the higher will be its basicity. Another useful trend is that basicity decreases as you go down a column of the periodic table. This is because the valence orbitals increase in size as one descends a column of the periodic table.

Complete answer:
Comparing the other two to ammonia, you will see that methylamine is a stronger base, whereas phenyl amine is very much weaker. All aliphatic primary amines are stronger bases than ammonia. Phenyl amine is typical of aromatic primary amines where the $ - N{H_2}$ group is attached directly to a benzene ring.
Amines are basic because they possess a pair of unshared electrons, which they can share with other atoms. These unshared electrons create an electron density around the nitrogen atom. The more basic the molecule the greater is the electron density.
The basicity of $N$ in ammonia, anilines attached to $N$ depends on the availability of the lone-pair of electrons on $N$. An aromatic ring is electron withdrawing and reduces the availability of the lone-pair of electrons on $N$, thereby reducing the basicity. In contrast, aliphatic carbons are electron donating, increasing the availability of the lone-pair on $N$ and thereby increasing the basicity.

Remember if there is an aromatic group in the molecule but not directly attached to the nitrogen of the amine. In this case, the amine will be more basic but not by much. For example, the\[p{K_a}\]of benzyl amine is $4.66$, slightly more basic than ammonia.