p-Block Elements (Groups 13 to 18)

In this chapter: general trends across the p-block, inert pair effect, anomalous behaviour of B, C, N, O and F, group-wise notes on groups 13 to 18, diborane, boric acid, allotropes of carbon, silicones and silicates, ammonia and nitric acid, oxoacids of phosphorus and sulphur, ozone and sulphuric acid, chlorine and interhalogen compounds, and the fluorides and oxides of xenon.

General trends

The p-block has the outer configuration ns²np¹⁻⁶ (helium, 1s², is placed in group 18). It contains metals, metalloids and non-metals; non-metallic character is highest at the top right and metallic character increases down each group.

• Oxidation states: the highest oxidation state equals the number of valence electrons (group number − 10). Down a group, the state two units lower becomes more stable. This inert pair effect arises because the ns² electrons of heavy elements are poorly shielded by d and f electrons and are reluctant to take part in bonding. So Tl⁺ is more stable than Tl³⁺, Pb²⁺ more stable than Pb⁴⁺ (PbO₂ is a strong oxidant), and Bi³⁺ more stable than Bi⁵⁺.
• Anomalous first member: B, C, N, O and F differ from the rest of their groups because of their small size, high electronegativity, high ionisation enthalpy and absence of d orbitals. Their maximum covalency is 4, and they form pπ-pπ multiple bonds (C=C, C≡C, N≡N, C=O, O=O) that heavier members form far less readily.
• Heavier members can use d orbitals to exceed a covalency of 4 (PCl₅, SF₆, IF₇) and can form dπ-pπ bonds, as in the oxides and oxoacids of P and S.

Group 13: B, Al, Ga, In, Tl

• Configuration ns²np¹; common oxidation state +3, with +1 increasingly stable down the group.
• Atomic radius increases down the group, except that Ga (135 pm) is smaller than Al (143 pm): the ten 3d electrons shield the nucleus poorly.
• Boron is a non-metal, very hard, and forms only covalent compounds. BX₃ molecules have only six electrons around boron, so they are electron deficient and act as Lewis acids (BF₃ + F⁻ → [BF₄]⁻). Aluminium chloride exists as the dimer Al₂Cl₆.
• Boric acid, B(OH)₃, is a weak monobasic Lewis acid: it accepts OH⁻ rather than releasing H⁺: B(OH)₃ + 2H₂O → [B(OH)₄]⁻ + H₃O⁺. It has a layered structure of planar BO₃ units held by hydrogen bonds. Borax is Na₂[B₄O₅(OH)₄]·8H₂O, used in the borax bead test for transition metal ions.

Diborane, B₂H₆

Diborane is a colourless, highly toxic gas that catches fire in air (B₂H₆ + 3O₂ → B₂O₃ + 3H₂O) and is hydrolysed by water to boric acid and hydrogen. Heating with ammonia finally gives borazine, B₃N₃H₆, called "inorganic benzene".

Each boron is roughly sp³ hybridised. The four terminal B-H bonds are ordinary two-centre two-electron bonds, and the two boron atoms and four terminal hydrogens lie in one plane. The two bridging hydrogens lie above and below this plane; each B-H-B bridge is a three-centre two-electron (3c-2e) bond, often called a banana bond. The bridge B-H bonds are longer than the terminal ones.

Group 14: C, Si, Ge, Sn, Pb

• Configuration ns²np²; oxidation states +4 and +2. The +2 state becomes more stable down the group: Sn²⁺ is a reducing agent (it tends to become Sn⁴⁺), while Pb⁴⁺ compounds are oxidising.
• Carbon is unique: it shows the greatest catenation (strong C-C bonds), forms pπ-pπ multiple bonds with itself and with O and N, and has a maximum covalency of 4.
• Allotropes of carbon: diamond (each C sp³, rigid three-dimensional network, hardest natural substance, non-conductor); graphite (each C sp², hexagonal layers held by weak van der Waals forces, conducts electricity, soft and slippery, used as a dry lubricant); fullerenes (C₆₀, sp² carbons in a cage of 20 six-membered and 12 five-membered rings).
• CO is poisonous because it binds haemoglobin to form carboxyhaemoglobin, about 300 times more stable than the oxygen-haemoglobin complex. CO₂ is linear (C is sp hybridised).
• SiO₂ is a three-dimensional network in which each Si is bonded tetrahedrally to four O atoms and each O to two Si atoms.

Silicones and silicates

Silicones are organosilicon polymers with the repeating unit -(R₂SiO)-. Methyl chloride reacts with silicon in the presence of copper at 573 K to give methyl-substituted chlorosilanes. Hydrolysis of dimethyldichlorosilane, (CH₃)₂SiCl₂, gives (CH₃)₂Si(OH)₂, which condenses into long chains; (CH₃)₃SiCl is added to cap the chain ends and control the chain length. Silicones repel water, resist heat and oxidation, are chemically inert and are good electrical insulators, so they are used in sealants, greases, electrical insulation and water-proofing of fabrics.

Silicates are built from SiO₄⁴⁻ tetrahedra. When the tetrahedra share corners with each other they form chains, rings, sheets or three-dimensional networks, and the negative charge is balanced by metal ions. Feldspars and zeolites are three-dimensional aluminosilicates; the zeolite ZSM-5 is used to convert alcohols directly into gasoline.

Group 15: N, P, As, Sb, Bi

• Configuration ns²np³, with a half-filled, extra stable p subshell. Oxidation states range from −3 to +5; the +3 state becomes more stable down the group.
• Nitrogen is anomalous: it exists as N₂ with a very strong triple bond (pπ-pπ), while phosphorus exists as P₄; nitrogen cannot expand its covalency beyond 4; and its single N-N bond is weak because of repulsion between lone pairs on the small atoms.
• Hydrides EH₃: basicity and thermal stability decrease down the group (NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃), while reducing character increases. The bond angle falls from NH₃ (107.8°) to PH₃ (93.6°), AsH₃ (91.8°) and SbH₃ (91.3°). NH₃ has an unusually high boiling point for its size because of hydrogen bonding.
• Phosphorus allotropes: white phosphorus (discrete tetrahedral P₄ molecules with 60° bond angles, strained and very reactive, glows in the dark, stored under water); red phosphorus (polymeric chains of P₄ units, much less reactive); black phosphorus (the most stable form, layered).
• PCl₅ is trigonal bipyramidal in the gas phase; its two axial bonds are longer, and weaker, than the three equatorial bonds. In the solid it exists as [PCl₄]⁺[PCl₆]⁻.

Ammonia

In the Haber process, N₂(g) + 3H₂(g) ⇌ 2NH₃(g), Δ_fH° = −46.1 kJ mol⁻¹. Following Le Chatelier's principle, high pressure (about 200 × 10⁵ Pa) favours ammonia; a temperature of about 700 K is a compromise between yield and rate, with iron oxide catalyst promoted by small amounts of K₂O and Al₂O₃. The NH₃ molecule is trigonal pyramidal, with three bond pairs and one lone pair on nitrogen. The lone pair makes it a Lewis base: with Cu²⁺ it forms the deep blue [Cu(NH₃)₄]²⁺.

Nitric acid

HNO₃ is a planar molecule and a strong oxidising agent. The product depends on concentration and on the metal:

Chromium and aluminium do not dissolve in concentrated HNO₃ because a protective oxide film forms (passivity). Concentrated HNO₃ oxidises non-metals: iodine to iodic acid, carbon to CO₂, sulphur to H₂SO₄ and phosphorus to H₃PO₄. In the brown ring test for nitrate, Fe²⁺ reduces nitrate to NO, which forms the brown complex [Fe(H₂O)₅(NO)]²⁺ at the junction with concentrated H₂SO₄.

Oxoacids of phosphorus

In all of them phosphorus is tetrahedral with one P=O bond. Only hydrogens attached to oxygen (P-OH) are ionisable; P-H hydrogens are not, but they give the acid reducing properties. On heating, H₃PO₃ disproportionates: 4H₃PO₃ → 3H₃PO₄ + PH₃.

Group 16: O, S, Se, Te, Po

• Configuration ns²np⁴. Oxygen shows mainly −2; S, Se and Te also show +2, +4 and +6, with +4 becoming more stable than +6 down the group.
• Oxygen is anomalous: small size and high electronegativity allow strong hydrogen bonding, so H₂O is a liquid while H₂S is a gas. Oxygen's electron gain enthalpy is less negative than that of sulphur because the added electron enters a small, crowded 2p subshell.
• Hydrides H₂E: acidic character increases (H₂O < H₂S < H₂Se < H₂Te) and thermal stability decreases down the group, as the H-E bond becomes longer and weaker.
• Ozone is a bent molecule (bond angle about 117°) with two equal O-O bonds of 128 pm, explained by resonance. It is a strong oxidising agent: it converts PbS to PbSO₄ and iodide to iodine.
• Sulphur exists as puckered S₈ rings (crown shape). Rhombic (α) sulphur is stable below 369 K; above it, monoclinic (β) sulphur is stable.
• SO₂ is bent; in water it acts as a reducing agent, decolourising acidified KMnO₄ and reducing Fe³⁺ to Fe²⁺.

Sulphuric acid and oxoacids of sulphur

In the contact process, sulphur or sulphide ore is burnt to SO₂, which is oxidised: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), Δ_rH° = −196.6 kJ mol⁻¹, over V₂O₅ at about 720 K and 2 bar. SO₃ is absorbed in H₂SO₄ to give oleum (H₂S₂O₇), which is diluted with water to the desired concentration.

• Dehydrating agent: it chars sugar (C₁₂H₂₂O₁₁ → 12C + 11H₂O) and removes water from many compounds.
• Oxidising agent (hot, concentrated): Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O; C + 2H₂SO₄ → CO₂ + 2SO₂ + 2H₂O.
• Oxoacids: sulphurous acid H₂SO₃ (S +4), sulphuric acid H₂SO₄ (+6), peroxodisulphuric acid H₂S₂O₈ (+6, contains an O-O peroxide link) and pyrosulphuric acid H₂S₂O₇ (+6, contains an S-O-S link). In each, sulphur is tetrahedral.

Group 17: F, Cl, Br, I

• Configuration ns²np⁵, one electron short of a noble gas. Halogens are strong oxidising agents; F₂ is the strongest, and oxidising power decreases down the group.
• Chlorine has the most negative electron gain enthalpy; fluorine's is less negative because of strong repulsion among electrons in its small 2p subshell.
• Bond dissociation enthalpy: Cl₂ > Br₂ > F₂ > I₂. The F-F bond is unexpectedly weak because of lone pair repulsion between the small atoms.
• Fluorine is anomalous: it shows only the −1 oxidation state, has no d orbitals, and HF is a liquid (hydrogen bonding) and a weak acid. Acid strength: HF < HCl < HBr < HI.
• Chlorine: MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O (laboratory); Deacon's process, 4HCl + O₂ → 2Cl₂ + 2H₂O (CuCl₂, 723 K). With cold dilute NaOH it gives NaCl + NaOCl; with hot concentrated NaOH, 3Cl₂ + 6NaOH → 5NaCl + NaClO₃ + 3H₂O. With dry slaked lime it gives bleaching powder: 2Ca(OH)₂ + 2Cl₂ → Ca(OCl)₂ + CaCl₂ + 2H₂O. Moist chlorine bleaches by oxidation (nascent oxygen), so the effect is permanent.
• Oxoacids of chlorine: HOCl (+1), HClO₂ (+3), HClO₃ (+5), HClO₄ (+7). Acid strength increases in the same order, HClO₄ being the strongest, because the conjugate base is increasingly stabilised by resonance over more oxygen atoms.

Interhalogen compounds

Halogens combine with each other to give compounds of types XX′, XX′₃, XX′₅ and XX′₇, where X is the larger, less electronegative halogen. They are covalent, diamagnetic and generally more reactive than the halogens themselves (except F₂), because the X-X′ bond is weaker than the X-X bond. Their shapes follow VSEPR theory:

Group 18: noble gases

Noble gases have closed-shell configurations (ns²np⁶; He 1s²), very high ionisation enthalpies and positive electron gain enthalpies, so they are very unreactive. Only xenon forms a well-known series of compounds. The first was made after it was noticed that the ionisation enthalpy of the O₂ molecule (1175 kJ mol⁻¹) is almost equal to that of xenon (1170 kJ mol⁻¹): since O₂⁺[PtF₆]⁻ was known, Xe⁺[PtF₆]⁻ was prepared.

Worked example: Use VSEPR theory to predict the shapes of XeF₄ and ClF₃.
Solution: Xe has 8 valence electrons; four are used in bonds with four F atoms, leaving 4 electrons = 2 lone pairs. Total 6 electron pairs, arranged octahedrally (sp³d²); the two lone pairs sit opposite each other, so XeF₄ is square planar. Cl has 7 valence electrons; three form bonds with F, leaving 4 electrons = 2 lone pairs. Total 5 pairs, trigonal bipyramidal (sp³d); both lone pairs take equatorial positions, so ClF₃ is bent T-shaped.

Xenon fluorides are strong fluorinating agents. XeF₆ is hydrolysed completely to XeO₃ (XeF₆ + 3H₂O → XeO₃ + 6HF); partial hydrolysis gives XeOF₄ and XeO₂F₂. Uses of noble gases: helium in weather balloons, in the gas mixture for deep-sea divers (it is much less soluble in blood than nitrogen) and in cryogenic work; neon in discharge tubes and signs; argon to provide an inert atmosphere in metallurgy and to fill electric bulbs.

Common mistakes: (1) Calling boric acid a protonic acid or a tribasic acid; it is a monobasic Lewis acid. (2) Counting P-H hydrogens when finding basicity; H₃PO₃ is dibasic and H₃PO₂ monobasic. (3) Taking F as having the most negative electron gain enthalpy; chlorine does. (4) Drawing XeF₄ as tetrahedral or ClF₃ as trigonal planar, forgetting lone pairs. (5) Explaining the smaller radius of Ga with the lanthanoid contraction; here it is the poor shielding by 3d electrons.

JEE and NEET focus

• Trends in radius, ionisation enthalpy, electron gain enthalpy and oxidation states, and the inert pair effect.
• Reasons for the anomalous behaviour of B, C, N, O and F.
• Structures: diborane, boric acid, silicones, PCl₅, oxoacids of P and S, ozone.
• Trends in hydrides of groups 15 and 16 and acid strength of hydrogen halides and oxoacids of chlorine.
• VSEPR shapes and hybridisation of interhalogens and xenon compounds.

Practice questions