Animated Solution for Chemistry - s and p-Block Elements: Based on the compounds of group 15 elements, the correct statement(s) is (are)

Imagine you are taking a walk down Group 15 of the periodic table, starting from the non-metallic Nitrogen at the top and ending at the heavy, metallic Bismuth at the bottom. This journey reveals a fascinating shift in physical and chemical properties. Let's break down the four statements given in the problem to understand the underlying chemistry.

As we move down any group in the periodic table, the atomic size increases, and the outermost electrons are held less tightly by the nucleus. This makes it easier for the atoms to lose electrons, which is the defining characteristic of a metal. Therefore, metallic character increases down the group.

Nitrogen ($extN$) is a pure non-metal. Non-metals typically form acidic oxides, meaning $ext{N}_{2}ext{O}_5$ is an acidic oxide. On the other hand, Bismuth ($extBi$) is a metal. Metals typically form basic oxides. Consequently, $ext{Bi}_{2}ext{O}_5$ is a basic oxide. This makes statement (A) absolutely correct: $ext{Bi}_{2}ext{O}_5$ is indeed more basic than $ext{N}_{2}ext{O}_5$.

Ionic Halides

The nature of a chemical bond depends heavily on the electronegativity difference between the bonding atoms. In $ext{NF}_3$, both Nitrogen and Fluorine are non-metals with high electronegativities. Because they are similar in nature, they share electrons, forming a covalent bond.

In contrast, $ext{BiF}_3$ consists of Bismuth (a metal) and Fluorine (a highly electronegative non-metal). This large difference in electronegativity leads to a significant transfer of electron density, giving the bond a strong ionic character. Therefore, $ext{NF}_3$ is much more covalent than $ext{BiF}_3$, making statement (B) correct.

Why does Ammonia ($ext{NH}_3$) behave so differently from Phosphine ($ext{PH}_3$)? The secret lies in hydrogen bonding.

Nitrogen is small and highly electronegative. When bonded to hydrogen, it pulls the shared electrons strongly towards itself, creating a significant partial negative charge on Nitrogen and a partial positive charge on Hydrogen. This allows $ext{NH}_3$ molecules to act like tiny magnets, sticking to each other through strong intermolecular hydrogen bonds.

Phosphorus is larger and less electronegative, so $ext{PH}_3$ cannot form these strong hydrogen bonds. Because it takes extra thermal energy to break the hydrogen bonds in Ammonia, $ext{NH}_3$ has a higher boiling point than $ext{PH}_3$. Thus, $ext{PH}_3$ boils at a lower temperature, making statement (C) correct.

N-N vs. P-P

You might intuitively think that because Nitrogen is more electronegative than Phosphorus, the $extN-N$ single bond should be stronger than the $extP-P$ single bond. However, there is a catch!

Nitrogen atoms are exceptionally small. When two Nitrogen atoms form a single bond, they have to get very close to each other. Picture two people trying to hold hands while wearing bulky backpacks. Those backpacks represent the lone pairs of electrons on each Nitrogen atom. Because the atoms are so close, these lone pairs repel each other violently. This severe interelectronic repulsion weakens the $extN-N$ bond.

Phosphorus atoms are larger, so the bond length is longer. The lone pairs are further apart, resulting in much less repulsion. As a result, the $extP-P$ single bond is actually stronger than the $extN-N$ single bond. This makes statement (D) incorrect.

By mastering these fundamental periodic trends, you can easily navigate complex inorganic chemistry problems!