The p-block elements occupy Groups 13 through 18 of the periodic table, filling their valence np subshells. Groups 13 through 15 are among the most reaction-dense sections of NEET inorganic chemistry, containing a rich variety of structural concepts, industrial processes, and property-based comparisons.

Group 13 — The Boron Family carries the outer electronic configuration $ns^{2}$ $np^{1}$, giving all members three valence electrons. Boron stands out as an anomaly: it is a nonmetal with a very small atomic radius and high electronegativity, forming exclusively covalent, electron-deficient compounds rather than the ionic compounds typical of heavier Group 13 members. This electron deficiency is the root cause of boron's Lewis acid character throughout its chemistry.

Diborane ($B_{2}H_{6}$) is the simplest and most important boron hydride. It cannot be predicted by classical Lewis structure theory — instead, it features two unusual 3-centre-2-electron (3c-2e) bonds, also called "banana bonds," where two boron atoms and one bridging hydrogen atom share just two electrons across all three. This electron-deficient bonding arises because $B_{2}H_{6}$ has only 12 valence electrons for what would classically require 14 for 7 two-centre bonds. Diborane has four terminal B-H bonds and two bridging B-H-B bonds; the molecule is often represented as $H_{2}B$(μ-H){2}$BH_{2}$. On hydrolysis: $B_{2}H_{6}$ + $6H{2}O → \2H_{3}BO_{3} + \6H_{2}$.

Borax ($Na_{2}B_{4}O_{7}$·$10H_{2}O$) is sodium tetraborate decahydrate. Its anion [$B_{4}O_{5}$(OH)_{4}]^{2-} is structurally notable for containing two trigonal planar$BO_{3}$units and two tetrahedral$BO_{4}$units — a mixed-unit composition frequently tested in NEET. The borax bead test exploits the ability of heated borax to form a glassy$B_{2}O_{3}$bead that dissolves metal oxides into characteristically coloured metaborates, enabling cation identification. Boric acid ($H_{3}BO_{3}$) is a weak monobasic Lewis acid. Unlike most acids, it does not donate $H^{+}$ directly. Instead, the electron-deficient boron atom (only 6 electrons in B(OH)_{3}) accepts $OH^{-}$ from water: $H_{3}BO_{3}$+$H_{2}O$→ [B(OH)_{4}]^{-} +$H^{+}$. The proton released originates from water, not from$H_{3}BO_{3}$ itself. In solid form, $H_{3}BO_{3}$molecules are connected into planar layers by O-H···O hydrogen bonds. Aluminium chloride ($AlCl_{3}$) dimerises to$Al_{2}Cl_{6}$because the monomer is electron-deficient; each Al achieves its octet by accepting a lone pair from a Cl of another$AlCl_{3}$unit via a coordinate bond. This dimer acts as a Lewis acid and is the classic Friedel-Crafts alkylation and acylation catalyst. The inert pair effect — the reluctance of the$ns^{2}$ electrons to participate in bonding — increases down Group 13, making Tl (+1) the most stable oxidation state for thallium.

Group 14 — The Carbon Family has the outer configuration $ns^{2}$ $np^{2}$. Carbon exhibits allotropy: diamond ($sp^{3}$, 3D tetrahedral network, hardest natural substance, electrical insulator), graphite ($sp^{2}$, planar hexagonal layers, delocalized π electrons conduct electricity, layers held by weak van der Waals forces making it a lubricant), and fullerene ($C_{60}$, spherical cage with 12 pentagonal and 20 hexagonal rings, $sp^{2}$ carbon). Carbon monoxide (CO) is a neutral ligand in coordination chemistry, coordinating through carbon to metal centres, and a potent poison that binds the $Fe^{2+}$ of haemoglobin approximately 200 times more strongly than $O_{2}$, forming carboxyhaemoglobin and preventing oxygen transport. Notably, magnesium is reactive enough to reduce C$O_{2}$: 2Mg + C$O_{2}$ → 2MgO + C, making C$O_{2}$ fire extinguishers lethal for Mg fires. Silicon forms silicones (linear or crosslinked –[$R_{2}SiO$]ₙ– polymers, water-repellent and thermally stable), silicates (built from $SiO_{4}^{4-}$ tetrahedral units linked in various dimensionalities), and zeolites (three-dimensional hydrated aluminosilicates used as molecular sieves and ion-exchange catalysts in water softening and petroleum cracking).

Group 15 — The Nitrogen Family has the outer configuration $ns^{2}$ $np^{3}$. The nitrogen molecule features an exceptionally strong N≡N triple bond (945 kJ/mol), making $N_{2}$ chemically inert at room temperature and requiring extreme industrial conditions to break it. The Haber process overcomes this: $N_{2}$ + $3H_{2}$⇌ 2N$H_{3}$, using a finely divided iron catalyst at 450°C and 200 atm. This is a compromise — lower temperatures would give better equilibrium yield (the forward reaction is exothermic) but unacceptably slow kinetics; higher pressure increases yield (4 mol reactants → 2 mol products) but increases equipment cost and safety risk. The Ostwald process converts N$H_{3}$to$HNO_{3}$in three steps: (1) catalytic oxidation of N$H_{3}$to NO over Pt-Rh gauze at 500°C; (2) uncatalysed oxidation of NO to N$O_{2}$at room temperature; (3) absorption of N$O_{2}$in water to give$HNO_{3}$, with NO recycled back to Step 2.

The five principal oxides of nitrogen span oxidation states +1 to +5: $N_{2}$O (+1, neutral, laughing gas), NO (+2, neutral, paramagnetic odd-electron molecule), N_{2}$$O_{3} (+3, acidic, anhydride of HN$O_{2}$), N$O_{2}$ (+4, acidic, brown, paramagnetic, dimerises to colourless N_{2}$$O_{4}), and N_{2}$$O_{5} (+5, acidic, anhydride of $HNO_{3}$). Phosphorus allotropes include white P ($P_{4}$ tetrahedra, 60° bond angles conferring high ring strain and reactivity, toxic, glows in dark due to slow oxidation, stored under water), red P (polymeric, much more stable, non-toxic), and black P (most thermodynamically stable, layered structure). $PCl_{5}$ adopts trigonal bipyramidal geometry with $sp^{3}$d hybridization, with 3 shorter equatorial bonds (~202 pm) and 2 longer axial bonds (~214 pm). For phosphorus oxoacids, the defining rule is that basicity equals the number of P-OH bonds, not the total hydrogen count: $H_{3}$P$O_{2}$ (1 P-OH, monobasic), $H_{3}$P$O_{3}$ (2 P-OH, dibasic), and $H_{3}$P$O_{4}$ (3 P-OH, tribasic). P-H bonds are non-ionizable and do not contribute to basicity.

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