Why is Primary Amine Stronger than Ammonia?

The premise behind the question 'Why is primary amine stronger than ammonia?' is fundamentally flawed. While it is true that the strength of a base can be influenced by the nature of the substituents attached to the nitrogen, the relative strength between primary amines and ammonia is not universally determined. The strength of a base, in the context of acidity and basicity, is typically quantified using the pKb value. The pKb value for a base is defined as the negative logarithm of its base dissociation constant, Kb. A lower pKb value indicates a stronger base.

Comparing Primary Amines and Ammonia

To gain a better understanding, let's look at some data:

Ammonia, NH3, in aqueous solution equilibrates as follows:

NH3(aq) H2O(l) rightleftharpoons NH4 (aq) HO-(aq)

pKb 4.76

On the other hand, a primary amine, such as methylamine, (CH3)3NH2, equilibrates as follows:

(CH3)3NH2(aq) H2O(l) rightleftharpoons (CH3)3NH3 (aq) HO-(aq)

pKb 3.44

Similarly, ethylamine, (CH3)2CHNH2, in aqueous solution equilibrates as follows:

(CH3)2CHNH2(aq) H2O(l) rightleftharpoons (CH3)2CHNH3 (aq) HO-(aq)

pKb 3.22

From these data, it is clear that primary amines, specifically methylamine and ethylamine, are indeed stronger bases than ammonia. This is primarily due to the electron-donating effect of the alkyl substituents in the primary amine.

Electronic Effects and Basicity

The relative basicity is greatly influenced by the nature of the substituents bonded to the nitrogen. In the case of primary amines, the alkyl groups (like methyl and ethyl) are electron-donating. This means that the alkyl groups donate electron density to the nitrogen atom, making it easier for the nitrogen to donate a proton and increasing the basicity of the amine.

In contrast, the widely used buffer Tryptamine (TRIS), which has an pKb value of approximately 8.0 at pH 8.0, is a much weaker base compared to both ammonia and primary amines. This lower basicity is associated with an electron-withdrawing group, which increases the net positive charge on the nitrogen and makes protonation more difficult.

For example, the hydroxymethyl group, -CH2OH, is an electron-withdrawing group. In a primary amine with a hydroxymethyl substituent, such as hydroxymethylamine, the pKb is significantly higher:

H2C-CH2NH2(aq) H2O(l) rightleftharpoons H2C-CH2NH3 (aq) HO-(aq)

pKb ≈ 9.0 (Note: This is an approximation, as the exact value can vary slightly)

This higher pKb indicates that hydroxymethylamine is a weaker base compared to primary amines with alkyl groups.

Influence of Substituents on Basicity

The electron-donating or electron-withdrawing nature of the substituent is crucial in determining the basicity. Electron-donating substituents, such as alkyl groups, increase the basicity by reducing the net positive charge associated with the ionization process, making it easier for the nitrogen to donate a proton. Electron-withdrawing substituents, on the other hand, increase the net positive charge, making ionization more difficult.

For instance:

Methylamine ((CH3)3NH2) is a stronger base than ammonia. Ethylamine ((CH3)2CHNH2) is a slightly stronger base than methylamine. Hydroxymethylamine (H2C-CH2NH2) is a weaker base than primary amines with alkyl groups.

Conclusion

In summary, the strength of a primary amine as a base compared to ammonia depends on the electron-donating or electron-withdrawing nature of the substituents. Alkyl groups are electron-donating and increase the basicity, while electron-withdrawing groups decrease it. Understanding these effects is crucial for predicting and manipulating the behavior of amines in different chemical contexts.