CBSE Class 12 Chemistry Chapter 7: The p-Block Elements - NCERT Solutions

The general characteristics of Group 15 elements (N, P, As, Sb, Bi) are as follows:

1. Electronic Configuration: All elements in Group 15 possess 5 valence electrons. Their general electronic configuration is $ns^2\; np^3$.
2. Oxidation States: Due to having 5 valence electrons, these elements typically show oxidation states of -3 (by gaining 3 electrons) and +3 or +5 (by losing or sharing valence electrons). The -3 oxidation state is more common for nitrogen and phosphorus. Nitrogen also exhibits -1 and -2 states. Down the group, the stability of the +5 oxidation state decreases, while the stability of the +3 oxidation state increases, primarily due to the inert pair effect.
3. Atomic Size: Atomic size increases as we move down the group from nitrogen to bismuth. This is because each subsequent element has an additional electron shell, increasing the distance of the valence electrons from the nucleus.
4. Ionisation Enthalpy: The first ionisation enthalpy generally decreases down the group. This trend is observed because the outermost electrons are further from the nucleus and are shielded by inner electrons, making them easier to remove. However, there might be minor irregularities due to changes in electronic configurations.
5. Electronegativity: Electronegativity decreases down the group. Nitrogen is the most electronegative element in this group. As the atomic size increases down the group, the attraction of the nucleus for the bonding electrons decreases, leading to lower electronegativity.

The reactivity of nitrogen differs significantly from that of phosphorus primarily due to the differences in their atomic structure and bonding capabilities:

1. Bond Strength: Nitrogen exists as a diatomic molecule (N₂) where two nitrogen atoms are linked by a very strong triple covalent bond ($N\; \backslash equiv\; N$). This triple bond has a very high bond dissociation enthalpy (946 kJ/mol), making it difficult to break and thus rendering nitrogen relatively inert under normal conditions.

2. pπ-pπ Bonding: Nitrogen is a small atom and can effectively form $p\backslash pi\; -\; p\backslash pi$ multiple bonds with itself. Phosphorus, being larger, cannot form stable $p\backslash pi\; -\; p\backslash pi$ bonds with itself effectively. While phosphorus can form $p\backslash pi\; -\; p\backslash pi$ bonds with other elements like oxygen, its own multiple bonding is less stable.

Consequently, phosphorus readily forms single bonds in its allotropes (like P₄), which are less stable and more reactive than the N₂ molecule. Therefore, phosphorus is considerably more reactive than nitrogen.

The chemical reactivity of Group 15 elements exhibits several trends:

1. Reactivity towards Hydrogen: All Group 15 elements react with hydrogen to form hydrides of the general formula $EH\_3$ (where E = N, P, As, Sb, Bi). The stability of these hydrides decreases significantly as we move down the group from $NH\_3$ to $BiH\_3$. For example, $NH\_3$ is a stable gas, while $BiH\_3$ is highly unstable and decomposes easily.
2. Reactivity towards Oxygen: Group 15 elements form oxides of the type $E\_2O\_3$ and $E\_2O\_5$. For example, nitrogen forms $N\_2O\_3$ and $N\_2O\_5$, while phosphorus forms $P\_4O\_6$ ($P\_2O\_3$) and $P\_4O\_\{10\}$ ($P\_2O\_5$). The acidic character of these oxides generally decreases down the group. Oxides in the higher oxidation state ($E\_2O\_5$) are typically more acidic than those in the lower oxidation state ($E\_2O\_3$).
3. Reactivity towards Halogens: These elements react with halogens to form halides of the type $EX\_3$ and $EX\_5$. However, nitrogen does not form $NX\_5$ because it lacks the d-orbitals in its valence shell, which are necessary to accommodate five bonding pairs of electrons. All trihalides, except for $NF\_3$ which is quite stable, are generally less stable than pentahalides where they can be formed.
4. Reactivity towards Metals: Group 15 elements react with metals to form binary compounds, typically exhibiting a -3 oxidation state. For instance, sodium nitride ($Na\_3N$) and calcium phosphide ($Ca\_3P\_2$) are examples of such compounds.

Ammonia ($NH\_3$) forms hydrogen bonds, whereas phosphine ($PH\_3$) does not significantly, due to the difference in electronegativity between nitrogen and phosphorus:

1. Electronegativity Difference: Nitrogen is highly electronegative (approx. 3.04 on the Pauling scale), while phosphorus is considerably less electronegative (approx. 2.19). This large difference in electronegativity between N and H in $NH\_3$ makes the N-H bonds highly polar. The nitrogen atom carries a significant partial negative charge ($\backslash delta^-$), and the hydrogen atoms carry a significant partial positive charge ($\backslash delta^+$).

2. Hydrogen Bonding: The highly polarized N-H bonds in $NH\_3$ allow the partially positive hydrogen atom of one $NH\_3$ molecule to be attracted strongly to the lone pair of electrons on the highly electronegative nitrogen atom of another $NH\_3$ molecule. This intermolecular attraction is known as hydrogen bonding.

3. Lack of Hydrogen Bonding in PH₃: In $PH\_3$, the P-H bonds are much less polar because phosphorus has a lower electronegativity. Consequently, the partial positive charge on the hydrogen atoms is very small, and the lone pair on the phosphorus atom is not sufficiently attractive to form strong intermolecular hydrogen bonds. Therefore, $PH\_3$ molecules are held together mainly by weaker van der Waals forces.