P- Block Elements

The p-block is the most fact-dense region of the periodic table in the NEET syllabus, contributing questions on trends, structures of oxoacids and oxides, and a long list of preparation reactions. Learning it as a set of patterns — with the first-member anomaly and the inert pair effect as your two master keys — is far more efficient than memorising each element separately.

What Defines the p-Block, and How Trends Are Organised

Elements in which the last electron enters a p-orbital of the outermost shell make up the p-block: groups 13 to 18. Their general valence configuration is ns²np¹⁻⁶ (helium is the exception, being 1s²). Since a p-subshell holds six electrons, each period contributes six p-block elements, and the block contains all the metalloids, most non-metals, and a few poor metals.

Two numbers control almost every trend:

• Group oxidation state = group number − 10 (so +3 for group 13, +4 for group 14, +5 for group 15, and so on). A lower state, two units below this, becomes progressively more stable down each group.
• Number of valence electrons decides the maximum covalency, but only for elements that have accessible d-orbitals. Second-period members (B, C, N, O, F) are restricted to a maximum covalency of four because they have no 2d orbitals; heavier congeners routinely show 5 and 6 (PCl₅, SF₆, IF₇).

Down a group, atomic radius and metallic character increase while ionisation enthalpy and electronegativity decrease, but the changes are not smooth. Two irregularities recur:

1. Post-transition contraction. Gallium is smaller than (or nearly equal to) aluminium, because the ten 3d electrons inserted before Ga screen the nucleus poorly. The same effect makes Ga's ionisation enthalpy higher than Al's.
2. Lanthanoid contraction. After the 4f series, Tl, Pb and Bi show unexpectedly high ionisation enthalpies and are reluctant to lose all their valence electrons.

The inert pair effect is the label for this reluctance: the ns² pair of heavy p-block elements is held tightly and does not participate in bonding, so the lower oxidation state dominates. Hence Tl(I) > Tl(III) in stability, Pb(II) > Pb(IV), Bi(III) > Bi(V). A direct consequence is oxidising power: PbO₂, BiO₃⁻ and Tl³⁺ are strong oxidants because they revert to the stable lower state.

The first member of every group behaves anomalously — small size, high electronegativity, high charge/radius ratio and no d-orbitals. This is why boron is a metalloid while the rest of group 13 are metals, why carbon shows catenation to an extreme degree, why nitrogen forms N≡N but phosphorus forms P₄, and why fluorine and oxygen are so different from their heavier relatives.

Groups 13 and 14: Boron, Aluminium, Carbon and Silicon Families

Group 13 (B, Al, Ga, In, Tl) has configuration ns²np¹. Boron is a hard, high-melting metalloid; the others are soft metals. The +3 state prevails at the top and +1 at the bottom.

Boron's compounds are dominated by electron deficiency — BX₃ has only six electrons around boron, so these are strong Lewis acids. Acid strength order is BF₃ < BCl₃ < BBr₃ < BI₃, the reverse of what electronegativity suggests, because in BF₃ the small fluorine 2p orbitals overlap efficiently with boron's empty 2p orbital (back-bonding), partially satisfying the deficiency.

Key group 13 chemistry to know:

• Borax, Na₂[B₄O₅(OH)₄]·8H₂O, gives an alkaline solution by hydrolysis; on heating it swells, then forms a clear glass of sodium metaborate and boric anhydride (the borax bead test).
• Orthoboric acid, H₃BO₃, is a layered solid held by hydrogen bonds; it is a weak monobasic Lewis acid — it accepts OH⁻ from water rather than donating H⁺: B(OH)₃ + H₂O → [B(OH)₄]⁻ + H⁺.
• Diborane, B₂H₆, has two three-centre-two-electron "banana" bonds bridging the borons; it burns with a huge exothermic enthalpy and is hydrolysed by water to boric acid.
• Aluminium dissolves in both acids and alkalies (amphoteric), giving [Al(H₂O)₆]³⁺ and [Al(OH)₄]⁻ respectively. Its oxide layer makes it corrosion-resistant.

Group 14 (C, Si, Ge, Sn, Pb) has ns²np². Catenation ability falls steeply C >> Si > Ge ≈ Sn >> Pb, tracking bond enthalpy (C–C is much stronger than Si–Si). Carbon alone forms stable multiple bonds with itself and with oxygen, so CO₂ is a discrete gas while SiO₂ is a three-dimensional covalent network of corner-sharing SiO₄ tetrahedra.

• Allotropes: diamond (sp³, tetrahedral network, hardest, non-conductor), graphite (sp² layers, delocalised electrons, conductor and lubricant), fullerenes (C₆₀, closed cage of 20 six-membered and 12 five-membered rings — the only pure, "molecular" allotrope).
• CO is a powerful ligand and a poison because it binds haemoglobin far more strongly than O₂. CO₂ is linear; its rising atmospheric concentration drives the greenhouse effect.
• Silicones are organosilicon polymers with (R₂SiO)ₙ chains, water-repellent and thermally stable. Silicates are built from SiO₄⁴⁻ units; zeolites are aluminosilicates used as catalysts (ZSM-5 for petroleum cracking) and ion exchangers.
• Halides: CCl₄ does not hydrolyse (no d-orbitals, no site for water to attack), whereas SiCl₄ hydrolyses readily. SnCl₂ is a reducing agent; PbCl₄ and PbO₂ are oxidising — the inert pair effect again.

Group 15 and Group 16: Nitrogen and Oxygen Families

Group 15 (N, P, As, Sb, Bi) — ns²np³ — shows states from −3 to +5. Nitrogen is unique: N₂ has a triple bond with enormous bond dissociation enthalpy (~941 kJ mol⁻¹), making dinitrogen almost inert at room temperature; phosphorus instead exists as P₄ tetrahedra with strained 60° angles, hence white phosphorus is highly reactive and glows in air (chemiluminescence).

• Ammonia is prepared industrially by the Haber process (Fe catalyst, ~200 atm, 700 K, moving mass by Le Chatelier). It is pyramidal, a good Lewis base and ligand, and gives characteristic precipitates/complexes with metal ions.
• Nitric acid comes from the Ostwald process: NH₃ → NO → NO₂ → HNO₃. Concentrated HNO₃ is a strong oxidant; with copper, dilute HNO₃ gives NO while concentrated gives NO₂. Zn behaves differently (very dilute acid can give N₂O/NH₄⁺). Passive metals (Cr, Al) and noble Au/Pt resist it.
• Phosphine, PH₃ — from hydrolysis of Ca₃P₂ or alkaline hydrolysis of white P — is less basic and much less hydrogen-bonded than NH₃, hence lower boiling point.
• Halides: PCl₃ hydrolyses to H₃PO₃; PCl₅ is trigonal bipyramidal in the gas phase but ionic, [PCl₄]⁺[PCl₆]⁻, in the solid.
• Oxoacids of phosphorus: count the P–H bonds to get basicity. H₃PO₄ (three OH, tribasic), H₃PO₃ (two OH + one P–H, dibasic, reducing), H₃PO₂ (one OH + two P–H, monobasic, strong reducing agent).

Group 16 (O, S, Se, Te, Po) — ns²np⁴. Oxygen shows −2 almost always (+2 only in OF₂); sulphur and beyond show +2, +4, +6. Hydride trend: H₂O has an anomalously high boiling point (hydrogen bonding); thereafter boiling point rises H₂S < H₂Se < H₂Te with size. Acid strength of hydrides increases down the group: H₂O < H₂S < H₂Se < H₂Te, because bond strength falls.

• Dioxygen is paramagnetic (two unpaired electrons in π* orbitals — molecular orbital theory). Ozone, O₃, is a bent resonance hybrid with equal bond lengths, a powerful oxidant that liberates I₂ from KI and turns Ag black; it shields Earth from UV and is destroyed catalytically by chlorine radicals from freons.
• Sulphur allotropes: rhombic (α, stable below 369 K) and monoclinic (β); both contain crown-shaped S₈ rings.
• SO₂ is bent, acts as both reductant and oxidant, and bleaches by reduction (temporary bleaching, colour returns).
• Sulphuric acid is made by the Contact process (SO₂ → SO₃ over V₂O₅, then absorption). It is a dehydrating agent, an oxidant when hot and concentrated, and a strong dibasic acid ionising in two steps.
• Structures: SF₆ octahedral and remarkably inert; SF₄ see-saw; H₂SO₄ tetrahedral around S with two S=O and two S–OH.

Group 17 (Halogens) and Group 18 (Noble Gases)

Halogens (F, Cl, Br, I, At) have ns²np⁵ and are one electron short of a stable octet, so they are the most reactive non-metals with high electron gain enthalpies. Note the anomaly: chlorine, not fluorine, has the most negative electron gain enthalpy, because fluorine's small size makes the incoming electron feel strong repulsion. Yet fluorine is the strongest oxidising agent in aqueous solution, owing to its low F–F bond dissociation enthalpy and the large hydration enthalpy of F⁻.

Key points:

• Bond dissociation enthalpy order: Cl₂ > Br₂ > F₂ > I₂ (F₂ is abnormally low — lone-pair repulsion in a tiny molecule).
• Hydrogen halides: HF is the weakest acid (strong H–F bond and hydrogen bonding); strength increases HF < HCl < HBr < HI.
• Oxidising power decreases F₂ > Cl₂ > Br₂ > I₂, so each halogen displaces the ones below it from their halides.
• Chlorine is prepared by oxidising HCl (MnO₂/KMnO₄) or by electrolysis of brine; it bleaches by oxidation (permanent). With cold dilute alkali it gives chloride + hypochlorite; with hot concentrated alkali, chloride + chlorate (3Cl₂ + 6NaOH(hot) → 5NaCl + NaClO₃ + 3H₂O) — the classic disproportionation pair examiners like to contrast with the cold-dilute case, where the same halogen disproportionates only as far as hypochlorite (Cl₂ + 2NaOH(cold, dilute) → NaCl + NaOCl + H₂O). Interhalogen compounds (ClF, BrF₃, ICl₅, IF₇) form between two different halogens, are always more reactive than either parent halogen (the X–X′ bond is weaker than X–X because the two atoms differ in size/electronegativity), and their shapes follow simple VSEPR counting once the central, larger, less electronegative halogen's lone pairs are included.

Group 18 (He, Ne, Ar, Kr, Xe, Rn) has a complete ns²np⁶ valence shell (1s² for helium), giving these elements the highest ionisation enthalpies in their periods and virtually no tendency to form compounds — hence the older name "inert gases." Only the heavier members, whose valence electrons are far enough from the nucleus to be somewhat available, form real compounds, and only with the most electronegative partners, fluorine and oxygen:

• Xenon fluorides: Xe + F₂ direct combination under controlled conditions gives XeF₂ (excess Xe), XeF₄ (1:5 ratio, higher pressure), or XeF₆ (large excess F₂, high pressure) depending on the Xe:F₂ ratio and conditions used.
• Shapes (VSEPR, treating lone pairs on Xe as included): XeF₂ is linear (3 lone pairs on Xe), XeF₄ is square planar (2 lone pairs), XeF₆ is distorted octahedral (1 lone pair).
• Hydrolysis of the fluorides gives oxo-compounds: XeF₂ hydrolyses slowly to Xe + HF + O₂; XeF₆ hydrolyses completely to XeO₃, a dangerously explosive solid.
• Xenon compounds are all strong oxidising and fluorinating agents themselves, and none is thermodynamically very stable — this entire branch of chemistry exists only because xenon's first ionisation enthalpy (1170 kJ mol⁻¹) is low enough, among the noble gases, for fluorine to pull electrons away.
• Helium, neon and argon form no confirmed stable compounds under ordinary conditions; helium's chemical inertness combined with its low density make it the safe choice for balloons and deep-sea diving gas mixtures (in place of flammable/narcotic-at-depth nitrogen), while argon's inertness is exploited to provide an unreactive atmosphere for arc welding and inside filament light bulbs.

Common Mistakes and Exam Traps

• Reading "group oxidation state = group number − 10" as the only stable state. It is simply the maximum theoretically accessible state; the inert pair effect often makes a lower state (two units below) the more stable one for the heaviest member of a group (Tl⁺, Pb²⁺, Bi³⁺), and NEET frequently asks which oxidation state is the more stable, not merely the maximum possible.
• Assuming BF₃ is the strongest Lewis acid among the boron trihalides because F is most electronegative. The order is actually reversed (BF₃ < BCl₃ < BBr₃ < BI₃) because F engages in back-bonding with boron's empty orbital, partially filling the electron deficiency that makes boron a Lewis acid in the first place — an inductive-effect assumption that fails once resonance/back-bonding is considered.
• Forgetting that CCl₄ does not hydrolyse while SiCl₄ does. Both are covalent tetrahalides of a group 14 element, but only silicon has energetically accessible d orbitals to accept the initial lone pair from an attacking water molecule; carbon's absence of low-lying d orbitals makes CCl₄ kinetically inert to hydrolysis despite being thermodynamically favourable.
• Mixing up which halogen has the most negative electron gain enthalpy versus which is the strongest oxidiser in solution. Chlorine (not fluorine) has the most negative electron gain enthalpy because fluorine's very small atom suffers strong inter-electron repulsion on gaining an electron, yet fluorine remains the strongest oxidising agent in aqueous solution once bond dissociation and hydration enthalpies are folded in — two different rankings that are frequently confused for the same thing.