Question

Pyridine is a base with $K_b$ value:

• A. $1.7 \times 10^{-9}$
• B. $2.3 \times 10^{-12}$
• C. $3.2 \times 10^{-6}$
• D. $3.8 \times 10^{-7}$

Hint:

Remember: Pyridine is a weak base with a $K_b$ around $10^{-9}$ — significantly less basic than aliphatic amines due to aromatic delocalization effects.

Answer 1

The Correct Option is A

- Pyridine is a heterocyclic aromatic organic compound with the molecular formula \ce{C5H5N}. It consists of a six-membered ring with five carbon atoms and one nitrogen atom.
- The nitrogen in pyridine is sp$^2$-hybridized and its lone pair of electrons is not involved in the aromatic $\pi$-system. This makes the lone pair available for protonation, hence making pyridine act as a Lewis base.
- The basicity of a compound is often expressed using the base dissociation constant ($K_b$), which measures the extent to which a base can accept a proton.
- Pyridine has a $K_b$ value of $1.7 \times 10^{-9}$, indicating it is a weak base. This low $K_b$ corresponds to a p$K_b$ of approximately 8.77.
- The low basicity compared to aliphatic amines is due to the electron-withdrawing effect of the aromatic ring, which reduces electron density on the nitrogen atom.
- Among the options, the correct $K_b$ value of pyridine is $1.7 \times 10^{-9}$.

Answer 2

The question gives four possible base dissociation constant values for pyridine and asks which one is correct. Since pyridine is a well known weak base, we can judge each value by how weak or strong a base it implies.

1. \(1.7 \times 10^{-9}\): A \(K_b\) this small corresponds to a pKb close to 8.8, placing pyridine as a distinctly weak base. This matches what we expect, since the nitrogen lone pair in pyridine sits in an sp2 orbital in the plane of the ring, away from the aromatic pi cloud, but the ring's electron-withdrawing character still pulls electron density off the nitrogen and limits how easily it grabs a proton.
2. \(2.3 \times 10^{-12}\): A value this tiny would place pyridine among extremely weak bases, weaker than even highly deactivated anilines. Pyridine is not that weak, since its lone pair is still available for protonation.
3. \(3.2 \times 10^{-6}\): This value is far too large. A \(K_b\) of this size would put pyridine close to simple aliphatic amines in basicity, but pyridine is known to be much weaker than aliphatic amines because its lone pair is held in a less basic sp2 orbital and the ring pulls on the nitrogen's electron density.
4. \(3.8 \times 10^{-7}\): This is also on the higher side for pyridine and would suggest a basicity closer to imidazole, which is not the case for the simple pyridine ring.

Weighing pyridine's known behavior against these four numbers, the correct order-of-magnitude fit is \(1.7 \times 10^{-9}\).

So the correct answer is \(1.7 \times 10^{-9}\).