Class 11 p-Block — The Boron and Carbon Families

Group 13 and Group 14 are where the Class 11 p-block chapter becomes real. Both groups begin with an element that refuses to behave like the rest of its family — boron and carbon — and both show a trend that reverses expectations halfway down. If you understand why each anomaly happens, you stop having to memorise a list of exceptions.

The two families at a glance

In both groups metallic character increases down the group, because the outer electrons get further from the nucleus and are lost more easily.

Trend 1 — atomic radius does not increase smoothly

Approximate atomic radii for Group 13 (in picometres) run B ≈ 85, Al ≈ 143, Ga ≈ 135, In ≈ 167, Tl ≈ 170. Read that again: gallium is slightly smaller than aluminium, even though it is one period lower.

The reason is the ten 3d electrons that appear between aluminium and gallium. The d electrons shield the outer electrons from the nucleus poorly, so the effective nuclear charge felt by gallium's outer electron is higher than expected, and the atom is pulled in tighter. This "d-block contraction" argument is a standard reasoning question in the board exam.

Trend 2 — the inert pair effect

The ns² pair of electrons becomes increasingly reluctant to take part in bonding, because the intervening d and f electrons shield the nucleus poorly and the ns pair is held more tightly. So:

• In Group 13, boron and aluminium show only +3, but Tl⁺ is more stable than Tl³⁺.
• In Group 14, carbon and silicon show +4, but Pb²⁺ is more stable than Pb⁴⁺.
• A direct consequence worth remembering: Sn²⁺ acts as a reducing agent (it wants to reach the more stable +4), while Pb⁴⁺ acts as an oxidising agent (it wants to drop to the more stable +2).

Why boron is the odd one out

Boron is anomalous for three connected reasons: it is very small, it has a high ionisation enthalpy, and it has no d orbitals in its valence shell.

• It never forms a simple B³⁺ ion — removing three electrons from such a small atom costs far too much energy — so all its compounds are covalent.
• Its maximum covalence is 4, not 6, because only the 2s and three 2p orbitals are available. Aluminium, with empty 3d orbitals, reaches a covalence of 6, which is why AlF₆³⁻ exists but BF₆³⁻ does not.
• In BF₃ and BCl₃ boron has only six electrons around it. These molecules are electron deficient and act as Lewis acids, accepting a lone pair: BF₃ + NH₃ → F₃B←NH₃.

The three boron compounds you must know

Borax, Na₂B₄O₇·10H₂O. Its solution is alkaline because it hydrolyses:

On strong heating borax loses its water, swells and then melts to a clear glassy bead:

That bead dissolves coloured metal oxides to give characteristic colours — the borax bead test.

Orthoboric acid, H₃BO₃. It is a soft, white, soapy-feeling solid with a layered structure held together by hydrogen bonds. It is a weak monobasic Lewis acid — and the word monobasic is the trap. It does not donate any of its three protons directly. Instead it accepts an OH⁻ from water:

On heating it loses water in stages: H₃BO₃ →(about 370 K) HBO₂ (metaboric acid) →(red heat) B₂O₃ (boric anhydride).

Diborane, B₂H₆. It has only 12 valence electrons for what looks like eight bonds, so it is electron deficient. Four hydrogens lie in the plane in ordinary two-centre two-electron bonds; the other two are bridging hydrogens held by three-centre two-electron "banana" bonds, in which one pair of electrons is shared over B–H–B. It burns and hydrolyses readily:

Aluminium is amphoteric

Aluminium dissolves in acid and in alkali, which is why aluminium vessels should not be used for either:

Worked example 1 — molar mass of borax and boric acid

Worked example 2 — percentage of boron in borax

Group 14 — why carbon is different

Carbon shows the same three causes as boron (small size, high ionisation enthalpy, no d orbitals) but the consequences are different, and two of them dominate the whole of organic chemistry.

• Catenation. Carbon links to itself far better than any other element in the group, because the C–C bond is unusually strong. Approximate single-bond enthalpies (kJ mol⁻¹): C–C ≈ 348, Si–Si ≈ 297, Ge–Ge ≈ 260, Sn–Sn ≈ 240. Lead effectively does not catenate. That falling order is the whole reason organic chemistry is carbon chemistry.
• pπ–pπ multiple bonding. Carbon's small 2p orbitals overlap sideways with oxygen's, so CO₂ is a discrete O=C=O molecule — a gas held together only by weak intermolecular forces. Silicon's 3p orbitals are too diffuse for good sideways overlap, so SiO₂ instead forms Si–O single bonds in every direction and becomes a giant covalent solid melting above 1900 K. Same formula type, completely different structure and properties.
• Maximum covalence 4. CCl₄ does not hydrolyse, because carbon has no d orbital to accept the incoming water's lone pair. SiCl₄ does: SiCl₄ + 4H₂O → Si(OH)₄ + 4HCl.

Allotropes of carbon, compared without a diagram

Both diamond and graphite are pure carbon. The difference in properties comes entirely from how the atoms are joined — a point board examiners like to test with "why does graphite conduct but diamond does not?".

Two industrial gases worth knowing

Water gas (CO + H₂) is made by passing steam over red-hot coke: C + H₂O → CO + H₂. Producer gas (CO + N₂) is made by passing air over hot coke. Carbon monoxide itself is highly poisonous because it binds to haemoglobin far more strongly than oxygen does, so it must only ever be handled in a properly ventilated laboratory under supervision.

Mistakes students repeat every year

Where this is examined

Always confirm the current syllabus for your board before you decide how deeply to revise a topic, since chapter coverage is revised from time to time.