Some p-Block Elements (Groups 13 and 14)NEET Chemistry · Class 11 · NCERT Chapter 12

Boron is the smallest element in Group 13 and has the highest charge density. Its compounds are largely covalent (e.g., BCl₃ is covalent, not ionic). Boron does not react with water or dilute acids under normal conditions, while Al reacts readily. Boron forms electron-deficient compounds (Lewis acids) and has a strong tendency to form multicentre bonds (like the 3c-2e bond in diborane). Boron also has an exceptionally high melting point (2300 °C) due to its covalent network structure, while other Group 13 metals are soft with low melting points. This is similar to the anomalous position of Li and Be in their respective groups.

The inert pair effect is the tendency of the outermost s-electrons (the ns² pair) to remain non-bonding in the heavier elements of p-block groups. As you go down a group, the ns² electrons become more stabilised due to poor shielding by inner d and f electrons, making them less available for bonding. In Group 13, Tl prefers +1 (not +3). In Group 14, Pb prefers +2 (not +4). In Group 15, Bi prefers +3. In Group 16, Po prefers +2. In Group 17, At and heavier halogens prefer lower oxidation states. The inert pair effect explains why heavier p-block elements show lower oxidation states.

Carbon has three main allotropes: (1) Diamond: each C is sp³ hybridised, forming a 3D tetrahedral network. It is the hardest natural substance, non-conductor, and has a very high melting point. (2) Graphite: each C is sp² hybridised, forming 2D hexagonal layers. The fourth electron is delocalised between layers, making graphite a good conductor of electricity and a lubricant (layers slide over each other). (3) Fullerenes (e.g., C₆₀, Buckminsterfullerene): C atoms arranged in pentagons and hexagons forming spherical cages. Each C is sp² hybridised. Fullerenes are semiconductors and have unique properties exploited in nanotechnology.

Diborane (B₂H₆) is the simplest boron hydride. Its structure is unusual because it has only 12 valence electrons for 7 bonds (2 B + 6 H). This is not enough for normal 2c-2e bonds everywhere. The molecule has two terminal B-H bonds on each boron (normal 2c-2e bonds) and two bridging B-H-B bonds. Each bridging bond is a 3-centre 2-electron (3c-2e) bond, where one electron pair is shared across three atoms (B-H-B). These are also called "banana bonds" or "bridge bonds". The four terminal H atoms and two boron atoms are coplanar; the two bridging H atoms are above and below this plane.

Carbon monoxide (CO) is a strong reducing agent: 2CO + O₂ → 2CO₂ and CO + metal oxide → metal + CO₂ (used in blast furnace to reduce iron ore). CO is extremely poisonous because it binds to haemoglobin (forming carboxyhaemoglobin) with an affinity 200 times greater than O₂. Carbon dioxide (CO₂) is generally not a strong reducing agent under normal conditions; it is already the oxidised product of carbon combustion. However, at very high temperatures, CO₂ can act as an oxidising agent for carbon: CO₂ + C → 2CO (Boudouard equilibrium). Both CO and CO₂ are acidic oxides, but CO₂ forms carbonic acid (H₂CO₃) with water while CO does not react with water directly.

Silicates are salts derived from silicic acid. The basic building unit is the SiO₄⁴⁻ tetrahedron. Different silicate types arise depending on how many corners (O atoms) are shared between tetrahedra: (1) Orthosilicate (island silicate): independent SiO₄⁴⁻ units, e.g., Mg₂SiO₄ (olivine). (2) Pyrosilicate: two SiO₄ units share one O, giving Si₂O₇⁶⁻. (3) Ring silicate: 3 or 6 SiO₄ units share 2 O each, giving Si₃O₉⁶⁻ or Si₆O₁₈¹²⁻, e.g., beryl. (4) Chain silicate: infinite chains, SiO₃²⁻ per unit, e.g., pyroxenes. (5) Double chain silicate: Si₄O₁₁⁶⁻ per unit, e.g., amphiboles. (6) Sheet silicate: 3 out of 4 O shared, Si₂O₅²⁻ per unit, e.g., micas, talc. (7) Framework silicate: all 4 O shared, SiO₂ (quartz, no net charge).