The p-Block Elements (Class 11) — NEET Notes

General characteristics of the p-block

The p-block runs from Group 13 to Group 18 — six columns, all defined by the rule that the differentiating electron enters a p-orbital. A p-subshell holds at most six electrons, which is why there are exactly six p-block groups; the valence shell configuration runs from ns²np¹ at boron to ns²np⁶ at neon. The block is unique because it holds the only non-metals on the table together with metalloids and metals — diversity that gives the p-block its complicated chemistry. Non-metallic character decreases down each group; the heaviest element in every p-block group is the most metallic. Non-metals form acidic or neutral oxides, metals form basic oxides, and the gradient between them is the cleanest way to trace metallic character.

The maximum oxidation state a p-block element can show equals its total number of valence electrons — Group 13 reaches +3, Group 14 reaches +4, Group 17 reaches +7. Alongside this group oxidation state, p-block elements also display states that differ from it by units of two, and which often become more stable as we move down the group. The first member of each p-block group differs sharply from its heavier congeners — small size, high electronegativity, no d-orbitals in the valence shell, and a unique ability to form pπ–pπ multiple bonds. These four factors together generate the anomalous behaviour of boron in Group 13 and carbon in Group 14.

The inert pair effect

Take a deliberate look at how the preferred oxidation state shifts as we walk down Groups 13, 14 and 15. Boron and aluminium are universally trivalent. Carbon and silicon are universally tetravalent. Nitrogen and phosphorus prefer +5 in their highest oxides. But by the time we reach the bottom of each group — thallium, lead, bismuth — the heavier element prefers an oxidation state two units lower: Tl⁺ over Tl³⁺, Pb²⁺ over Pb⁴⁺, Bi³⁺ over Bi⁵⁺. The ns² pair of outermost electrons stops participating in bonding. This reluctance is the inert pair effect.

The cause is the poor shielding of the ns electrons by the intervening d and f electrons in heavier members. Across the long sixth period the 4f and 5d electrons screen the 6s pair very inefficiently, the effective nuclear charge on the 6s pair grows, and those electrons are held so tightly that promoting them into the valence state — the energy required to break the pair — exceeds what subsequent bond formation can recover. NIOS adds a second contributor: poorer orbital overlap in the larger atoms means each bond formed in the higher oxidation state is weaker than in the lighter homologue, so the energetic incentive to promote those s electrons shrinks.

Group 13 — the boron family

Group 13 contains boron, aluminium, gallium, indium and thallium (plus the synthetic nihonium). The valence configuration is ns²np¹. Boron is a typical non-metal — hard, black, refractory, with an unusually high melting point because its solid is a crystalline covalent network. Aluminium is a familiar silvery metal; on a weight-for-weight basis its electrical conductivity is twice that of copper. Gallium is the curiosity — its melting point is so low (303 K) that it liquefies on a warm palm, yet its boiling point (2676 K) makes it useful for high-temperature thermometers. Density rises down the group from 2.35 g cm⁻³ for B to 11.85 g cm⁻³ for Tl.

Atomic radius increases from B (88 pm) to Al (143 pm) as expected, but then there is a deviation — gallium (135 pm) is smaller than aluminium because gallium sits after the first-row d-block. The ten 3d electrons added between Al and Ga screen the nuclear charge poorly, so Ga's outer electron experiences a stronger effective pull and contracts. The same logic explains the irregular electronegativity trend: it falls from B (2.0) to Al (1.5), then climbs marginally back to 1.8 at Tl.

Down the group, the nature of the oxides changes systematically. B₂O₃ is acidic. Al₂O₃ and Ga₂O₃ are amphoteric. In₂O₃ and Tl₂O₃ are basic. This is the classic gradient of metallic character read off the oxide row.

The anomaly of boron

Boron stands apart from the rest of Group 13 for the same four reasons that every first-row p-block element stands apart: tiny atomic size, high electronegativity, high ionisation enthalpy, and no d-orbitals in the valence shell. The first three together drive the sum of B's first three ionisation enthalpies so high that B³⁺ cation formation is energetically prohibitive — boron is locked into covalent bonding. The fourth point is the most testable: with only 2s and three 2p orbitals available, boron's maximum covalence is four. It forms tetrahedral [BF₄]⁻ by accepting a fluoride into its empty p-orbital, but it cannot form octahedral [BF₆]³⁻ the way aluminium does, because that would require six bonds and there are only four orbitals.

This restricted covalence is what forces boron's most distinctive compounds — diborane, the electron-deficient hydride; boric acid, which acts as a Lewis acid rather than a Brønsted acid; and borax, with its hidden [B₄O₅(OH)₄]²⁻ ring — into shapes the heavier elements never adopt.

Borax, boric acid, and diborane

Three boron compounds dominate the NEET footprint of this chapter: borax (a cheap industrial mineral), orthoboric acid (a weak antiseptic), and diborane (the textbook example of an electron-deficient hydride). All three are interconvertible in the laboratory along a single route, and NEET keeps testing the structure of each one.

Borax — Na₂B₄O₇·10H₂O

Borax is the most important compound of boron — a white crystalline solid found in nature at Puga Valley (Ladakh) and Sambhar Lake (Rajasthan). The traditional formula Na₂B₄O₇·10H₂O hides a more accurate structural formula Na₂[B₄O₅(OH)₄]·8H₂O, in which the anion is a tetranuclear ring of two trigonal BO₃ and two tetrahedral B(OH)₄ units fused through bridging oxygens. Aqueous borax is alkaline because hydrolysis releases NaOH alongside orthoboric acid:

Na₂B₄O₇ + 7 H₂O → 2 NaOH + 4 H₃BO₃

Hydrolysis of borax — NEET 2018 / 2019 motif

On heating, borax loses water and swells up; further heating turns it into a transparent liquid that solidifies into a glassy bead of sodium metaborate plus boric anhydride. The borax bead test exploits this: transition-metal oxides fused with borax on a platinum loop give characteristically coloured metaborates — cobalt gives a blue Co(BO₂)₂ bead, chromium gives green, manganese gives violet.

Orthoboric acid — H₃BO₃

Orthoboric acid is a white crystalline solid with a soapy touch, sparingly soluble in cold water but freely so in hot water. It has a striking layer structure: planar BO₃ triangles are linked through hydrogen bonds in two-dimensional sheets — which is why it is polymeric (NEET MCQ 11.34 of old NCERT exercise). Despite its name, it is not a Brønsted acid. It does not donate a proton; instead, it pulls a hydroxide from water and releases H₃O⁺, making it a weak monobasic Lewis acid:

B(OH)₃ + 2 H₂O ⇌ [B(OH)₄]⁻ + H₃O⁺

Boric acid acts as a Lewis acid by accepting OH⁻

On heating above 370 K, H₃BO₃ dehydrates stepwise — first to metaboric acid HBO₂, then to boric oxide B₂O₃.

Diborane — B₂H₆

The simplest boron hydride is diborane — a colourless, highly toxic gas with a boiling point of 180 K that catches fire spontaneously on exposure to air. It is prepared in the laboratory by oxidising sodium borohydride with iodine, and industrially by treating BF₃ with NaH at 450 K.

Diborane is the canonical electron-deficient hydride. Twelve valence electrons are available, but eight ordinary 2-centre-2-electron bonds would need fourteen. The molecule resolves this by using four normal terminal B–H bonds and two 3-centre-2-electron bridge bonds (the famous "banana bonds"). Each boron uses sp³ hybrid orbitals. Two terminal hydrogens and the two borons lie in one plane; the two bridging hydrogens sit above and below. NEET 2022 Q.54 attacked exactly this geometry and offered "sp²" as a wrong-answer trap.

Diborane reacts with Lewis bases to cleave into BH₃ adducts (B₂H₆ + 2 NMe₃ → 2 BH₃·NMe₃). With ammonia it gives borazine B₃N₃H₆ — "inorganic benzene" — a six-membered ring of alternating BH and NH units. Boron's important hydridoborate anion [BH₄]⁻ is the working species inside NaBH₄ and LiBH₄, both standard reducing agents in organic synthesis.

Aluminium and AlCl₃

Aluminium is the most abundant metal in the earth's crust (8.3% by mass) — third overall after oxygen and silicon. Its principal minerals are bauxite (Al₂O₃·2H₂O) and cryolite (Na₃AlF₆). Aluminium combines low density (2.70 g cm⁻³) with high tensile strength and conducts electricity at twice the per-gram efficiency of copper — properties that make it dominant in transmission cables, aerospace alloys, packaging, and cookware. The metal is electropositive (E°_Al³⁺/Al = −1.66 V) and prefers a single +3 oxidation state across all its compounds.

Aluminium is also amphoteric — a property NEET routinely turns into one-line statements. It dissolves in dilute HCl liberating H₂: 2 Al + 6 HCl → 2 AlCl₃ + 3 H₂. It dissolves in aqueous NaOH too, forming the tetrahedral aluminate anion: 2 Al + 2 NaOH + 6 H₂O → 2 Na[Al(OH)₄] + 3 H₂. Concentrated HNO₃ passivates aluminium by depositing a protective oxide film — which is why conc. HNO₃ can be transported in aluminium containers.

Anhydrous AlCl₃ exists as the chlorine-bridged dimer Al₂Cl₆. Each Al atom forms three normal Al–Cl bonds and accepts a lone pair from a chlorine on the neighbouring AlCl₃ unit to reach a tetrahedral octet. This dimerisation is the structural answer to AlCl₃'s electron deficiency, and it is why anhydrous AlCl₃ "smokes" in air — atmospheric moisture hydrolyses the dimer, releasing HCl gas. In acidified aqueous solution, Al³⁺ exists as the octahedral hexaaqua complex [Al(H₂O)₆]³⁺ with sp³d² hybridisation at aluminium, employing the 3d orbitals that boron lacks. Potash alum, KAl(SO₄)₂·12H₂O, is one of many "alum" double salts.

Uses of boron and aluminium compounds

Boron fibres are used in bullet-proof vests and lightweight composite materials for aircraft. The ¹⁰B isotope absorbs neutrons effectively, so metal borides serve as protective shields and control rods in nuclear reactors. Borax and boric acid go into heat-resistant glasses such as Pyrex, glass-wool, and fibreglass. Borax also functions as a soldering flux, a glaze for earthenware, and a constituent of medicinal soaps. Dilute aqueous boric acid is a mild antiseptic — useful for eyewashes and skin preparations.

Aluminium and its alloys are used in everything from aircraft fuselages to soft-drink cans. Cryolite is the molten medium for industrial Al extraction by electrolysis. Anhydrous AlCl₃ is the workhorse Lewis-acid catalyst in Friedel–Crafts reactions. Potash alum is used in water purification (settles colloidal impurities) and as a styptic.

Group 14 — the carbon family

Group 14 contains carbon, silicon, germanium, tin and lead (with the synthetic flerovium below). The valence configuration is ns²np². Carbon and silicon are non-metals, germanium is a metalloid, and tin and lead are soft metals with low melting points. Carbon ranks seventeenth in crustal abundance but is the central element of all organic chemistry and of every living organism; its three natural isotopes ¹²C, ¹³C and ¹⁴C (the radioactive one used for dating, half-life 5,770 years) cover most analytical needs. Silicon is the second most abundant element on earth (27.7% by mass), present as silica and silicates throughout the crust.

Covalent radius increases sharply from C (77 pm) to Si (118 pm), then increases only slightly from Si through Pb. The reason is the same one we met in Group 13 — filled d and f orbitals in the heavier members screen the nucleus poorly. First ionisation enthalpies fall down the group, but not smoothly: Sn to Pb shows a small increase because the 4f electrons in Pb screen the 6s pair so weakly that the effective nuclear charge climbs. Electronegativity is highest for carbon (2.5) and roughly constant near 1.8 for Si through Pb.

Down the group the dioxides become progressively less acidic: CO₂, SiO₂ and GeO₂ are acidic; SnO₂ and PbO₂ are amphoteric. Among monoxides CO is unique — neutral; GeO is distinctly acidic; SnO and PbO are amphoteric.

The anomaly of carbon and catenation

Carbon, like boron in Group 13, differs sharply from its heavier congeners. Small atomic radius (77 pm), high electronegativity (2.5), high ionisation enthalpy and absence of d-orbitals in its valence shell together force maximum covalence 4. Carbon cannot form [CF₆]²⁻, but silicon forms [SiF₆]²⁻ readily — Si uses its 3d orbitals; C has none. The same restriction explains why CCl₄ resists hydrolysis (carbon can't accept a water lone pair) while SiCl₄ hydrolyses rapidly through a five-coordinate transition state on silicon.

Carbon's signature property is catenation — the formation of chains and rings of like atoms held by covalent bonds. The reason is the unusually high C–C bond enthalpy (about 348 kJ mol⁻¹). As we move down Group 14, atomic size increases, bond enthalpy falls, and catenation weakens dramatically:

Carbon is also the only Group 14 element that forms strong pπ–pπ multiple bonds — both with itself (C=C, C≡C in ethene and ethyne) and with other small electronegative atoms (C=O in carbonyls, C=N in imines, C≡N in nitriles). Heavier Group 14 elements have atomic orbitals that are too large and too diffuse for effective sideways p-overlap, so silicon and germanium analogues of these multiple bonds are exceptionally rare. Combined, catenation and pπ–pπ bonding make carbon the most allotrope-rich element in the periodic table.

Allotropes of carbon

Carbon exists in several crystalline and many amorphous forms. The three crystalline allotropes — diamond, graphite and fullerenes — have appeared on NEET in some form almost every year. Diamond is the hardest natural substance; graphite is the softest crystalline carbon; fullerenes are cage-like molecules discovered as late as 1985 by Kroto, Smalley and Curl, who collected the Nobel Prize for it in 1996. NEET 2023 Q.68 was a direct List-I/List-II match on these three forms plus coke.

Graphite is thermodynamically the most stable allotrope of carbon (ΔH_f° = 0 by definition). Diamond, with ΔH_f° = +1.90 kJ mol⁻¹, is metastable and would in principle convert to graphite — except the activation energy at room conditions is impossibly high. Fullerene, at +38.1 kJ mol⁻¹, is the least stable of the three but still kinetically inert. Amorphous "carbons" — coke, charcoal, carbon black — are impure forms of graphite or fullerene. Coke is used as a reducing agent in metallurgy (NEET 2023 Q.68 matched this). Activated charcoal adsorbs poisonous gases in water filters and air conditioners; carbon black goes into automobile tyres as a filler.

Carbon monoxide and carbon dioxide

The two industrial oxides of carbon — CO and CO₂ — sit on the boundary between organic and inorganic chemistry. Carbon monoxide is produced by incomplete combustion or by passing steam over hot coke at 473–1273 K to give "water gas" (CO + H₂), also called synthesis gas — a key industrial fuel and feedstock. When air rather than steam is used, the product is "producer gas" (CO + N₂). CO is colourless, odourless and almost water-insoluble. It is a powerful reducing agent and reduces most metal oxides — including Fe₂O₃ to Fe and ZnO to Zn — which is the basis of much extractive metallurgy in the blast furnace.

CO has the structure :C≡O: with one σ and two π bonds. The lone pair on carbon makes CO a powerful ligand — it forms metal carbonyls like Ni(CO)₄ and Fe(CO)₅ with transition metals. Its toxicity comes from the same lone pair: CO binds to the iron of haemoglobin to form carboxyhaemoglobin, a complex about 300 times more stable than oxyhaemoglobin. The result is that the blood can no longer carry oxygen and tissues asphyxiate. NEET 2020 Q.172 tested exactly this — the correct trap-detection was that carboxyhaemoglobin is more, not less, stable than oxyhaemoglobin.

Carbon dioxide is generated by complete combustion of carbon or hydrocarbons, by acid-on-limestone in the laboratory (CaCO₃ + 2 HCl → CaCl₂ + H₂O + CO₂), and commercially by heating limestone. CO₂ is linear (O=C=O) with sp-hybridised carbon. Two sp orbitals on C form σ bonds with O; the remaining two p electrons participate in pπ–pπ bonding with oxygen's p-orbitals. The C–O bond lengths are equal (115 pm), and the molecule has no dipole moment.

CO₂ dissolves slightly in water to give the dibasic carbonic acid H₂CO₃, which is in equilibrium with HCO₃⁻ and CO₃²⁻. The carbonate / bicarbonate buffer keeps blood pH between 7.26 and 7.42. Atmospheric CO₂ sits at about 0.03% by volume, removed by photosynthesis and added by combustion of fossil fuels — the balance is what drives the greenhouse effect and modern global warming. Solid CO₂ ("dry ice") is used as a refrigerant; gaseous CO₂ carbonates soft drinks; it is heavy and non-supporter of combustion, so it works as a fire extinguisher.

Silicon — silica, silicates, silicones, zeolites

Silicon and oxygen between them make up roughly 95% of the earth's crust. The full hierarchy of silicon–oxygen compounds — from simple SiO₂ through silicates, silicones and zeolites — is built from a single repeat unit: the SiO₄⁴⁻ tetrahedron, in which each silicon is covalently bonded to four oxygens in a tetrahedral arrangement.

Silicon dioxide — SiO₂

SiO₂ occurs as quartz, cristobalite and tridymite — interconvertible crystalline forms — and as amorphous opal. The structure is a three-dimensional covalent network in which every Si is bonded to four O atoms and every O bridges two Si atoms, giving eight-membered rings of alternating Si and O. The very high Si–O bond enthalpy makes silica almost inert: it resists most acids and metals even at red heat. The two reagents that do attack it are HF and aqueous NaOH:

SiO₂ + 4 HF → SiF₄ + 2 H₂O  ·  SiO₂ + 2 NaOH → Na₂SiO₃ + H₂O

Why HF etches glass; why hot caustic dissolves silica

Quartz is piezoelectric — applied pressure generates a voltage — which is why quartz crystals time wristwatches and stabilise radio oscillators. Silica gel is used as a desiccant and as a chromatographic support. Kieselguhr, an amorphous silica, packs filtration plants.

Silicates — orthosilicate to tectosilicate

When SiO₄ tetrahedra link to each other by sharing oxygens at corners, the resulting silicate anions classify into five neat structural families, distinguished by how many oxygens each tetrahedron shares with neighbours. NEET treats this classification as gettable factual knowledge.

Silicones — (R₂SiO)_n polymers

Silicones are synthetic organosilicon polymers built around the (R₂SiO)_n repeat unit. The starting material is an alkyl- or aryl-substituted silicon chloride R_nSiCl_(4−n), obtained by reacting methyl chloride with silicon at 573 K in the presence of copper catalyst. Hydrolysis of (CH₃)₂SiCl₂ followed by condensation polymerisation gives a straight-chain silicone; adding (CH₃)₃SiCl caps the ends and controls the chain length. Because silicones are wrapped in non-polar alkyl groups, they are hydrophobic, thermally stable, dielectric, and chemically inert. Their applications run from sealants and high-temperature greases to electrical insulators, water-proofing for fabrics, and biocompatible implants in surgery and cosmetic procedures.

Zeolites — open-framework aluminosilicates

Zeolites are tectosilicates in which some Si⁴⁺ centres are replaced by Al³⁺. Each substitution drops the framework charge by one unit; the deficit is balanced by exchangeable cations — Na⁺, K⁺, Ca²⁺ — sitting in the channels of the lattice. The open framework gives zeolites three industrial uses NEET likes to test:

• Catalysts: ZSM-5 cracks hydrocarbons and converts methanol to gasoline. NEET 2020 Q.158 tested this directly.
• Ion exchangers: Permutit water softeners use sodium zeolites; the Na⁺ in the framework is exchanged for the Ca²⁺/Mg²⁺ in hard water.
• Molecular sieves: the channels admit small molecules selectively, used to dry solvents and separate gas mixtures.

Common natural zeolites include natrolite Na₂Al₂Si₃O₁₀·2H₂O, heulandite Ca[Al₂Si₇O₁₈]·6H₂O and chabazite. Synthetic zeolites such as ZSM-5 dominate modern catalytic chemistry — petrochemical cracking towers run on aluminosilicate beds.