Main Group Elements: VSEPR, Compounds, Boranes and Silicates, Allotropes, Industrial Chemistry and Acid–Base Concepts

Valence-shell electron-pair repulsion places the bonding and lone pairs of a central atom as far apart as possible. The steric number (bonded atoms + lone pairs) fixes the electron geometry: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral, 7 pentagonal bipyramidal. Repulsions fall in the order lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, so lone pairs compress bond angles (CH₄ 109.5°, NH₃ 107°, H₂O 104.5°). In a trigonal bipyramid lone pairs go equatorial (fewer 90° contacts) and the more electronegative ligands go axial; in an octahedron two lone pairs go trans. Multiple bonds count as one region but repel more strongly.

Hydrides are saline (ionic, with H⁻: LiH, CaH₂), covalent (molecular: CH₄, NH₃, H₂O, B₂H₆) or metallic/interstitial (PdHₓ). Electron-deficient diborane B₂H₆ has four terminal B–H bonds and two three-centre two-electron B–H–B bridges. Halides of the p-block hydrolyse when the central atom has a vacant low-lying orbital to accept water: SiCl₄ and PCl₅ hydrolyse readily, while CCl₄ and SF₆ are kinetically inert because carbon has no d orbitals and sulfur is sterically shielded. Oxides become more acidic across a period and with rising oxidation state (Na₂O basic, Al₂O₃ amphoteric, SiO₂ weakly acidic, P₄O₁₀ and SO₃ strongly acidic), and more basic down a group.

Oxoacids. The basicity of a phosphorus oxoacid equals its number of P–OH groups, not its number of hydrogens: H₃PO₄ is tribasic, H₃PO₃ (one P–H) dibasic and H₃PO₂ (two P–H) monobasic and a good reducing agent. For oxoacids of one element, acidity rises with the number of oxygens not bearing hydrogen (Pauling’s rule): HClO < HClO₂ < HClO₃ < HClO₄. Sulfur forms H₂SO₃, H₂SO₄, H₂S₂O₇ (disulfuric, oleum) and the peroxoacids H₂SO₅ and H₂S₂O₈, in which sulfur is still +6 and the extra oxidising power sits in the O–O bond. Nitrides range from ionic Li₃N and Mg₃N₂ (which give NH₃ with water) through covalent BN and Si₃N₄ to interstitial TiN; sulfides from ionic Na₂S to the layered MoS₂ and the cage molecules P₄S₃ and S₄N₄.

Wade’s rules (polyhedral skeletal electron pair theory) count the electrons holding a borane cage together. Each B–H unit gives 2 skeletal electrons, each extra H 1, a C–H unit 3, and a negative charge adds its electrons. With n vertices, n + 1 skeletal pairs means a closo cage (a closed deltahedron: B₆H₆²⁻ octahedron, B₁₂H₁₂²⁻ and C₂B₁₀H₁₂ icosahedra); n + 2 a nido cage (one vertex missing: B₅H₉, B₁₀H₁₄); n + 3 an arachno cage (two missing: B₄H₁₀, B₅H₁₁). The styx numbers describe the same molecules in terms of 3c–2e B–H–B bridges (s), 3c–2e BBB bonds (t), 2c–2e B–B bonds (y) and extra BH₂ groups (x).

The isolobal analogy (Hoffmann) says two fragments are isolobal when their frontier orbitals have the same number, symmetry, approximate energy and occupancy. CH₃ and Mn(CO)₅ each offer one singly occupied orbital pointing outward; CH₂ and Fe(CO)₄ offer two frontier orbitals with two electrons; CH and Co(CO)₃ offer three orbitals with three electrons. In the same sense a B–H vertex is isolobal with C–H⁺, which is why replacing BH⁻ by CH turns a borane anion into a neutral carborane. So Mn₂(CO)₁₀ is the metal analogue of ethane, and Co(CO)₃ units can replace CH vertices in tetrahedrane-like clusters — the bridge between borane, hydrocarbon and metal-cluster chemistry.

Silicates are built from SiO₄ tetrahedra sharing corners. The number of shared (bridging) oxygens per silicon sets the formula: 0 in orthosilicates SiO₄⁴⁻ (olivine), 1 in pyrosilicates Si₂O₇⁶⁻, 2 in rings (Si₆O₁₈¹²⁻ in beryl) and single chains (SiO₃)ₙ²ⁿ⁻ (pyroxenes), 2.5 in double chains Si₄O₁₁⁶⁻ (amphiboles, asbestos), 3 in sheets Si₂O₅²⁻ (micas, clays, talc) and 4 in three-dimensional frameworks SiO₂ (quartz), where replacing Si by Al gives aluminosilicate frameworks (feldspars, zeolites). Silicones (R₂SiO)ₙ come from hydrolysis and condensation of R₂SiCl₂; R₃SiCl caps chains and RSiCl₃ cross-links them, giving oils, rubbers and resins that are water-repellent and thermally stable. Boron nitride is isoelectronic with carbon: hexagonal BN is layered like graphite but an insulator (the B–N polarity localises the π electrons), and cubic BN, like diamond, is superhard. Borazine B₃N₃H₆, "inorganic benzene", is planar and isoelectronic with benzene but its polar B–N bonds make it add HCl or water readily. Phosphazenes (NPCl₂)ₙ, from PCl₅ and NH₄Cl, form rings (n = 3, 4) and high polymers with alternating P–N bonds; substituting Cl by OR gives stable elastomers.

Carbon: diamond (sp³, three-dimensional, hardest, insulating), graphite (sp² layers 335 pm apart, conducting in-plane, thermodynamically stable at ambient conditions), fullerenes (C₆₀ with 12 pentagons and 20 hexagons), nanotubes and graphene. Phosphorus: white P₄ (tetrahedral, P–P–P angle 60°, strained, very reactive, glows in air and is stored under water), red (polymeric chains of P₄ units, less reactive) and black (puckered layers, the most stable, a semiconductor). Sulfur: rhombic α-S₈ (crown rings, stable below 95.5 °C), monoclinic β-S₈ (stable from 95.5 °C to the melting point) and plastic sulfur (long helical chains from quenching the melt).

Industrial synthesis is an exercise in equilibrium and kinetics. Ammonia (Haber–Bosch): N₂ + 3H₂ ⇌ 2NH₃ is exothermic (ΔH° ≈ −92 kJ/mol) with Δn = −2, so high pressure (about 150–300 atm) raises the yield, and low temperature would too, but at a useless rate; a compromise near 400–450 °C is used with a promoted iron catalyst (K₂O, Al₂O₃), and NH₃ is condensed out and unreacted gas recycled. Sulfuric acid (contact process): S or sulfide ores are burned to SO₂; 2SO₂ + O₂ ⇌ 2SO₃ (exothermic) is run over V₂O₅ at about 400–450 °C; SO₃ is absorbed in concentrated H₂SO₄ to give oleum, H₂S₂O₇, which is diluted, because SO₃ with water directly forms a fine acid mist. Nitric acid (Ostwald process): 4NH₃ + 5O₂ → 4NO + 6H₂O over a Pt–Rh gauze at about 850 °C with very short contact time; NO is oxidised to NO₂, which is absorbed in water, 3NO₂ + H₂O → 2HNO₃ + NO, and the NO is recycled.

Noble gases. Xenon, with the lowest ionisation energy of the stable noble gases, forms compounds with the most electronegative elements. Heating Xe with F₂ gives XeF₂ (linear), XeF₄ (square planar) or XeF₆ (distorted octahedral) depending on the ratio and conditions — shapes that VSEPR predicts from three, two and one lone pairs. XeF₆ hydrolyses to XeO₃, a trigonal pyramidal molecule that is dangerously explosive when dry; XeF₂ is a clean fluorinating agent. Interhalogens XY, XY₃, XY₅ and XY₇ (X the larger halogen; IF₇ is the only XY₇) have shapes from VSEPR (ClF₃ T-shaped, BrF₅ square pyramidal, IF₇ pentagonal bipyramidal) and are more reactive than the parent halogens, except F₂, because the X–Y bond is weaker and polar; ClF₃ and BrF₃ are powerful fluorinating agents. Polyhalide ions such as I₃⁻ (linear) form when a halide coordinates a halogen molecule.

A Brønsted acid donates a proton and a base accepts one; every acid has a conjugate base, and the stronger the acid the weaker its conjugate base. In water the strongest acid that can exist is H₃O⁺ (the levelling effect), so HClO₄, HCl and HNO₃ all look equally strong; a less basic solvent such as acetic acid differentiates them. A Lewis acid accepts an electron pair and a base donates one, which covers metal ions, BF₃ and AlCl₃. Lewis acidity of the boron trihalides runs BF₃ < BCl₃ < BBr₃, the reverse of electronegativity, because the strong B–F π back-donation into boron’s empty p orbital must be lost when BF₃ pyramidalises to accept a base.

Pearson’s hard and soft acids and bases sorts Lewis acids and bases by polarisability. Hard acids are small, highly charged and weakly polarisable (H⁺, Li⁺, Na⁺, Mg²⁺, Al³⁺, Ti⁴⁺, Fe³⁺, BF₃); soft acids are large, low-charged and polarisable (Cu⁺, Ag⁺, Au⁺, Hg²⁺, Pd²⁺, Pt²⁺, zero-valent metals). Hard bases: F⁻, OH⁻, H₂O, NH₃, O²⁻, RCOO⁻; soft bases: I⁻, S²⁻, RS⁻, CN⁻, CO, PR₃, H⁻, SCN⁻ bound through S; borderline: Fe²⁺, Co²⁺, Ni²⁺, Cu²⁺, Zn²⁺, Br⁻, pyridine, N₃⁻. The principle — hard prefers hard, soft prefers soft — explains why Al and Ti occur as oxides and Hg, Pb and Cu as sulfides, why Ag⁺ binds I⁻ far more strongly than F⁻, and why thiocyanate binds hard metals through N and soft ones through S. Acid–base catalysis speeds reactions by protonating a substrate (ester hydrolysis, keto–enol tautomerism, dehydration) or by deprotonating it (aldol reactions); Lewis acids such as AlCl₃ catalyse Friedel–Crafts reactions by generating the electrophile.