CBSE Class 12 Chemistry Chapter 14: 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: These elements typically exhibit oxidation states of -3, +3, and +5. The -3 oxidation state arises from gaining three electrons to complete their octet. Nitrogen, being small, can achieve this more readily. The +3 and +5 states involve the valence electrons. The stability of the +5 oxidation state decreases down the group, while the stability of the +3 oxidation state increases due to the 'inert pair effect'. Nitrogen also shows -1 and -2 oxidation states.
3. Atomic Size: Atomic radius 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 ionization 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.
5. Electronegativity: Electronegativity decreases down the group. Nitrogen is the most electronegative element in this group. This decrease is primarily due to the increase in atomic size, which reduces the attraction of the nucleus for bonding electrons.

The difference in reactivity between nitrogen and phosphorus stems from their distinct structural and bonding properties:

• Bond Strength: Nitrogen exists as a diatomic molecule (N₂) with a strong triple covalent bond (N≡N). This triple bond has very high bond dissociation enthalpy, making it difficult to break and thus rendering nitrogen gas relatively unreactive under normal conditions.
• pπ-pπ Bonding: Nitrogen, being a small atom, can effectively form stable pπ-pπ multiple bonds with itself. Phosphorus, being larger, cannot form such stable pπ-pπ bonds with itself. Instead, it forms single bonds, leading to molecules like P₄ (tetrahedral structure) where P-P single bonds are present. These single bonds are easier to break than the triple bond in N₂.

Consequently, phosphorus is significantly more reactive than nitrogen.

The chemical reactivity of Group 15 elements shows distinct trends when reacting with other elements:

1. Reactivity towards Hydrogen: Group 15 elements react with hydrogen to form hydrides of the general formula $EH\_3$ (where E = N, P, As, Sb, Bi). Examples include $NH\_3$, $PH\_3$, $AsH\_3$, $SbH\_3$, and $BiH\_3$. The stability of these hydrides decreases significantly as we move down the group from $NH\_3$ to $BiH\_3$.
2. Reactivity towards Oxygen: These elements form oxides of the type $E\_2O\_3$ and $E\_2O\_5$. For example, nitrogen forms $N\_2O\_3$ and $N\_2O\_5$. The acidic character of these oxides generally decreases down the group. Oxides where the element is in a higher oxidation state are typically more acidic than those where it is in a lower oxidation state.
3. Reactivity towards Halogens: Group 15 elements react with halogens to form halides of the type $EX\_3$ and $EX\_5$. However, nitrogen does not form $NX\_5$ compounds because it lacks the d-orbitals in its valence shell, limiting its covalency to a maximum of four. Most of these halides are stable compounds.
4. Reactivity towards Metals: Group 15 elements react with metals to form binary compounds. In these compounds, the Group 15 element typically exhibits a -3 oxidation state, acting as the more electronegative element. For example, sodium nitride ($Na\_3N$) is formed when nitrogen reacts with sodium.

The ability of ammonia ($NH\_3$) to form hydrogen bonds, while phosphine ($PH\_3$) does not, is primarily due to the difference in electronegativity between nitrogen and phosphorus:

• Electronegativity Difference: Nitrogen is significantly more electronegative than phosphorus. This higher electronegativity causes nitrogen to attract the bonding electrons in the N-H bond more strongly, creating a substantial partial positive charge ($\backslash delta+$) on the hydrogen atoms and a partial negative charge ($\backslash delta-$) on the nitrogen atom in $NH\_3$.
• Hydrogen Bonding in $NH\_3$: The highly polarized N-H bonds in ammonia allow the partially positive hydrogen atoms of one $NH\_3$ molecule to form strong electrostatic attractions (hydrogen bonds) with the lone pair of electrons on the partially negative nitrogen atom of another $NH\_3$ molecule.
• Lack of Hydrogen Bonding in $PH\_3$: Phosphorus is less electronegative than nitrogen. Consequently, the P-H bonds in phosphine ($PH\_3$) are much less polarized. The partial positive charge on the hydrogen atoms and the partial negative charge on the phosphorus atom are too small to facilitate significant hydrogen bonding between $PH\_3$ molecules.

Therefore, ammonia exhibits intermolecular hydrogen bonding, leading to a higher boiling point and solubility compared to phosphine.