The p-Block ElementsNEET Chemistry · Class 12 · NCERT Chapter 13

The inert pair effect is the tendency of the outermost s-electrons (the ns² pair) to remain non-bonding in heavier elements. In Group 15, the common oxidation states are +5 and +3. As you go down the group from N to Bi, the +3 state becomes increasingly stable while the +5 state becomes less common. This happens because the 6s² electrons in Bi (and 5s² in Sb) are held tightly by poor shielding from d and f electrons below them, making them harder to use for bonding. So bismuth mainly shows +3, and nitrogen mainly shows +5 (or -3). The trend: N (+5 most common) > P > As > Sb > Bi (+3 most stable).

Nitrogen cannot form pentahalides (such as NF5 or NCl5) because it does not have d orbitals available in its valence shell. Nitrogen is in period 2 and uses only s and p orbitals (maximum 4 bonds, octet rule). Phosphorus is in period 3 and has vacant 3d orbitals. These 3d orbitals can be used to expand the octet, allowing P to form five bonds. PCl5 and PF5 exist with sp3d hybridization of phosphorus. Similarly, N cannot form more than 4 bonds, but P can form up to 5. This is one of the key anomalous behaviors of nitrogen compared to the rest of Group 15.

White phosphorus: discrete P4 molecules with a tetrahedral structure. The bond angle is only 60°, which causes ring strain, making it highly reactive. It catches fire spontaneously in air above 30°C. It is waxy, soft, white/yellow, and extremely toxic. It glows in the dark (chemiluminescence). Red phosphorus: a polymeric form where P4 units are linked by covalent bonds into chains. No discrete small molecules. Much more stable than white P. Does not ignite spontaneously. Used in safety match boxes. Non-toxic. Black phosphorus: the most stable allotrope, made at high pressure. Has a layered structure like graphite. Least reactive. A semiconductor.

PCl5 structure: phosphorus is sp3d hybridized (5 bond pairs, 0 lone pairs). The geometry is trigonal bipyramidal. There are two types of bonds: three equatorial P-Cl bonds (bond angle 120°) and two axial P-Cl bonds (bond angle 90° with equatorial). The axial bonds are longer and weaker than equatorial bonds because axial bonding involves p orbitals with more repulsion. In the gas phase, PCl5 dissociates partly into PCl3 and Cl2. SF6 structure: sulfur is sp3d2 hybridized (6 bond pairs, 0 lone pairs). The geometry is perfectly octahedral (bond angle 90°). All six S-F bonds are equivalent. SF6 is chemically very inert because the large SF6 molecule prevents any reagent from approaching the central S atom (steric protection).

The boiling point order among hydrogen halides is: HF (19.5°C) >> HI (-35.4°C) > HBr (-66.8°C) > HCl (-85.1°C). HF has an anomalously high boiling point because fluorine is the most electronegative element. The H-F bond is highly polar, and HF molecules form strong intermolecular hydrogen bonds (F-H...F). These hydrogen bonds require more energy to break, giving HF a much higher boiling point. HCl, HBr, and HI cannot form significant hydrogen bonds (Cl, Br, I are not electronegative enough to form strong H-bonds). Among HCl, HBr, and HI, the boiling point increases with molecular mass due to increasing van der Waals forces.

H2SO3 (sulphurous acid): S is in +4 oxidation state. It is a weak diprotic acid formed when SO2 dissolves in water. It is a good reducing agent. H2SO4 (sulphuric acid): S is in +6 oxidation state. Strong diprotic acid. Made industrially by the contact process. Acts as an acid, dehydrating agent, and oxidising agent (concentrated). It reacts with water vigorously (exothermic). H2S2O7 (oleum or pyrosulphuric acid/disulphuric acid): formed when SO3 is dissolved in concentrated H2SO4. S is in +6. It is more strongly acidic and oxidising than H2SO4. Used industrially when making H2SO4 via the contact process. On adding water: H2S2O7 + H2O → 2H2SO4. Other oxoacids include H2S2O8 (peroxodisulphuric acid, S in +6 with peroxy linkage).

The ozone layer (stratosphere, 15-40 km altitude) absorbs harmful UV-B and UV-C radiation from the sun before it reaches Earth. Without this shield, UV radiation would cause skin cancer, cataracts, and damage to ecosystems. Ozone absorbs UV: O3 + UV → O2 + O. The O atom recombines: O + O2 → O3. This cycle continuously absorbs energy. Destruction by CFCs: chlorofluorocarbons (Freons) release Cl atoms in the stratosphere: CF2Cl2 + UV → CF2Cl + Cl. The Cl atom acts as a catalyst: Cl + O3 → ClO + O2, then ClO + O → Cl + O2. Net: O3 + O → 2O2. One Cl atom can destroy thousands of O3 molecules. Nitrogen oxides (NOx from supersonic jets) also destroy ozone by a similar catalytic mechanism.

Interhalogen compounds are compounds formed between two different halogens. The general formula is XYn where X is the heavier halogen and n = 1, 3, 5, or 7. They are more reactive than individual halogens (the X-Y bond is weaker than X-X or Y-Y because it is polar). Types: XY (ClF, BrF, BrCl, ICl, IBr) - linear. XY3 (ClF3 T-shaped sp3d hybridization 2 lone pairs, BrF3 T-shaped) - T-shaped geometry. XY5 (ClF5, BrF5 - square pyramidal sp3d2 hybridization with 1 lone pair, IF5) - square pyramidal. XY7 (IF7 only - pentagonal bipyramidal sp3d3 hybridization with 0 lone pairs). All interhalogen compounds are reactive oxidizing agents. ICl (iodine monochloride) can be used as a source of I+ ions in electrophilic iodination reactions.