The basicity of amines refers to their ability to accept a proton (H⁺) because of the lone pair of electrons on the nitrogen atom. In simple terms, it tells us how “strong” a base the amine is. Here’s a detailed explanation:
1. Why Amines Are Basic
- The nitrogen atom in amines has a lone pair of electrons.
- This lone pair can accept a proton, forming an ammonium ion.
- The general reaction is: amine + H⁺ → ammonium ion.
So, the more readily the nitrogen donates its lone pair to H⁺, the stronger the base.
2. Factors Affecting the Basicity of Amines
a) Alkyl Substituents (Inductive Effect):
- Alkyl groups are electron-donating.
- They push electron density towards nitrogen, making the lone pair more available to accept a proton.
- This increases basicity in general:
- Secondary amine > Primary amine > Ammonia in water.
- Tertiary amines are slightly less basic than secondary in water because of steric hindrance and solvation effects.
b) Aromatic Amines (Anilines):
- In aniline, the lone pair on nitrogen is partially delocalized into the benzene ring.
- This reduces the availability of the lone pair for protonation.
- Therefore, aromatic amines are less basic than aliphatic amines.
c) Solvent Effect:
- In water, solvation stabilizes the ammonium ion.
- Sterically hindered amines (like tertiary amines) are less well solvated, slightly reducing their basicity compared to secondary or primary amines.
3. Relative Basicity
- Aliphatic amines (RNH₂, R₂NH) > Ammonia > Aromatic amines (ArNH₂) in aqueous solution.
- Tertiary amines are less basic than secondary amines in water, despite having more alkyl groups, because the bulky groups hinder solvation of the protonated form.
Summary
- Amines are basic due to the lone pair on nitrogen.
- Electron-donating groups increase basicity.
- Delocalization or steric hindrance decreases basicity.
- Solvent plays a role in stabilizing the protonated amine.