Class 11 Chemistry — Chapter 304: Elements of Group 13, 14 and 15

Key concepts: Elements of Group 13, 14 and 15 (Class 11 Chemistry)

This chapter covers the first three families of the p-block: group 13 (ns²np¹, the boron family), group 14 (ns²np², the carbon family) and group 15 (ns²np³, the nitrogen family). In each family the head element — boron, carbon and nitrogen — behaves abnormally because it is small, highly electronegative and has no d orbitals in its valence shell, which caps its covalency at four and pushes it towards pπ-pπ multiple bonding. Below the head element the trends settle down and two new ideas appear: the inert pair effect, which makes the lower oxidation state more stable down group 13 and 14, and the availability of d orbitals, which lets silicon and phosphorus expand their octets. The chapter then works through the important compounds — borax, boric acid, diborane, aluminium compounds, the allotropes of carbon, silica, silicates and silicones, ammonia, nitric acid, the oxides of nitrogen, the allotropes of phosphorus, phosphine, phosphorus halides and the oxoacids of phosphorus.

Group 13 trends

Boron is a metalloid while Al, Ga, In and Tl are metals. The group oxidation state is +3, and atomic radius does not rise smoothly — gallium is marginally smaller than aluminium because the poorly shielding 3d electrons added before it raise the effective nuclear charge. Gallium also has an exceptionally low melting point of about 303 K.

Inert pair effect

Going down groups 13 and 14, the ns² pair of electrons becomes progressively reluctant to take part in bonding because the intervening d and f electrons shield the nucleus poorly and the s electrons are held more tightly. Hence Tl⁺ is more stable than Tl³⁺ and Pb²⁺ more stable than Pb⁴⁺, which makes Tl³⁺ and Pb⁴⁺ strong oxidising agents.

Anomalous behaviour of boron

Boron is far smaller and more electronegative than the rest of its group and its valence shell offers only four orbitals, so its maximum covalency is four. It never forms a simple B³⁺ ion, its trihalides are electron-deficient Lewis acids, and it forms [BF₄]⁻ but never [BF₆]³⁻, unlike aluminium which uses 3d orbitals to give [AlF₆]³⁻.

Borax

Borax, Na₂B₄O₇·10H₂O, is the commonest boron mineral and gives an alkaline solution on hydrolysis. Heated strongly it swells, loses its water and finally melts to a clear glassy bead of sodium metaborate and boric anhydride; the B₂O₃ in this bead combines with coloured metal oxides to give the characteristic colours of the borax bead test.

Boric acid

H₃BO₃ is a soft, soapy, white solid whose planar B(OH)₃ units are linked into layers by hydrogen bonds. It is a weak monobasic Lewis acid — it does not release its own protons but accepts OH⁻ from water to give [B(OH)₄]⁻ and H₃O⁺. On heating it gives metaboric acid HBO₂ and finally boric oxide B₂O₃.

Diborane

B₂H₆ has only 12 valence electrons for eight atoms and is therefore electron deficient. Four hydrogens form ordinary terminal B–H bonds while two bridge the borons through three-centre two-electron 'banana' bonds; the bridge bonds are longer than the terminal ones. Diborane burns in oxygen with a huge release of energy and is hydrolysed by water to boric acid and hydrogen.

Aluminium and its compounds

Aluminium is amphoteric: it dissolves in acids to give Al³⁺ and in alkali to give the aluminate ion [Al(OH)₄]⁻ with hydrogen, and concentrated nitric acid renders it passive through an oxide film. Anhydrous AlCl₃ exists as the dimer Al₂Cl₆, acts as a Lewis acid catalyst in Friedel-Crafts reactions, and potash alum KAl(SO₄)₂·12H₂O is used to purify water.

Catenation and covalency in group 14

The strength of the E–E bond decides catenation, which falls steeply as C >> Si > Ge ≈ Sn, with lead showing practically none. Carbon is limited to a covalency of four, while silicon and the heavier members can use vacant d orbitals to reach six, which is why SiCl₄ is readily hydrolysed by water but CCl₄ is not, and why [SiF₆]²⁻ exists but [CF₆]²⁻ does not.

Diamond and graphite

In diamond every carbon is sp³ hybridised and bonded tetrahedrally to four others in a rigid three-dimensional network, making it the hardest natural substance and a non-conductor. In graphite sp² carbons form hexagonal sheets with one delocalised electron each, so it conducts electricity, and the widely spaced layers slide over one another, making it a lubricant.

Fullerenes

Buckminsterfullerene C₆₀ is a closed cage of twenty six-membered and twelve five-membered rings shaped like a football, made by heating graphite in an electric arc in an inert atmosphere. All its carbons are sp² hybridised, and because it has no dangling bonds it is the only pure and smooth allotrope of carbon.

Oxides of carbon

CO is a neutral, sparingly soluble and highly poisonous gas prepared by dehydrating formic acid with concentrated sulphuric acid; it binds haemoglobin about 300 times more strongly than oxygen does. CO₂ is an acidic oxide of discrete O=C=O molecules, exists as dry ice when solid, and is the main greenhouse gas released by burning fuels.

Silica and silicates

SiO₂ is a giant covalent network in which each Si is bonded to four O atoms and each O to two Si atoms, which is why it is hard and very high melting while CO₂ is a gas. Silicates are built from SiO₄⁴⁻ tetrahedra sharing corners — sharing two corners gives chains, three gives sheets and all four gives three-dimensional frameworks such as quartz.

Silicones

Silicones are organosilicon polymers with an alternating Si–O–Si backbone, made by hydrolysing alkyl chlorosilanes such as (CH₃)₂SiCl₂ and letting the resulting silanols condense. R₃SiCl caps a chain, R₂SiCl₂ gives linear chains and RSiCl₃ gives cross-linked networks; the outward-facing alkyl groups make silicones water repellent and heat stable.

Anomalous behaviour of nitrogen

Nitrogen is small, highly electronegative and lacks d orbitals, so it forms strong pπ-pπ multiple bonds and exists as the triply bonded N₂ gas, whereas phosphorus forms single bonds in the P₄ tetrahedron. Its maximum covalency is four, so it forms no pentahalide, and N–N single bonds are weak because of lone pair repulsion, which is why phosphorus catenates better than nitrogen.

Ammonia and nitric acid

NH₃ is made by the Haber process over promoted iron at about 700 K and 200 atm; it is pyramidal, sp³ hybridised with a bond angle of about 107°, and acts as a Lewis base forming complexes such as the deep blue [Cu(NH₃)₄]²⁺. HNO₃ is made by the Ostwald process in which ammonia is oxidised over a Pt/Rh gauze to NO, then to NO₂, which disproportionates in water.

Oxides of nitrogen

Nitrogen forms N₂O and NO, both neutral, and N₂O₃, NO₂, N₂O₄ and N₂O₅, all acidic. N₂O₃ is the anhydride of nitrous acid and N₂O₅ of nitric acid, while brown NO₂ exists in equilibrium with colourless N₂O₄. Nitric acid never gives hydrogen with metals: dilute acid gives NO with copper and concentrated acid gives NO₂.

Allotropes of phosphorus

White phosphorus is a waxy, poisonous solid of discrete P₄ tetrahedra whose 60° bond angles are badly strained, so it glows in the dark, ignites in air near 303 K and must be stored under water. Red phosphorus is a polymeric chain of linked P₄ units and is far less reactive, while black phosphorus is the most stable form and conducts electricity.

Phosphine and phosphorus halides

PH₃ is prepared by heating white phosphorus with concentrated NaOH in an inert atmosphere; it is a weaker base and far less soluble than ammonia, and impure samples catch fire spontaneously. PCl₃ is a pyramidal liquid hydrolysed to H₃PO₃, while PCl₅ is trigonal bipyramidal as a gas but exists as [PCl₄]⁺[PCl₆]⁻ in the solid state.

Oxoacids of phosphorus

Basicity equals the number of P–OH groups, not the number of hydrogens in the formula. H₃PO₂ is monobasic with two P–H bonds, H₃PO₃ is dibasic with one P–H bond and both are good reducing agents, whereas H₃PO₄ is tribasic, has no P–H bond and is not a reducing agent; H₄P₂O₇ is tetrabasic.