Q.Give reasons for the following:

Diamond (sp3, rigid 3-D network, no free electrons) is hard and non-conducting; graphite (sp2, layered, one free electron per C) is soft and conducting. Carbon's superior catenation ability (stronger, more stable C-C bonds than Si-Si) explains why alkanes vastly outnumber silanes.

1. Diamond vs graphite: In diamond, every carbon atom is sp3-hybridised and forms four strong sigma (C-C) bonds to four neighbouring carbon atoms, arranged tetrahedrally, extending in all three dimensions into a giant, rigid covalent network. Since all four valence electrons of each carbon are used in strong, localised sigma bonds, there are no free/mobile electrons — this is why diamond is extremely hard (a rigid, strongly-bonded 3-D lattice) and a non-conductor of electricity (no charge carriers). In graphite, each carbon is sp2-hybridised and forms only three sigma bonds to three neighbouring carbons, all lying in the same plane, building up flat hexagonal (honeycomb) layers. Successive layers are held together only by weak van der Waals forces, so the layers can easily slide over one another — this is why graphite is soft and slippery (used as a lubricant/in pencils). Each carbon's fourth valence electron occupies an unhybridised p-orbital perpendicular to the layer; these p-electrons are delocalised over the whole layer, free to move — this delocalised electron cloud makes graphite a good conductor of electricity (along the layers).
2. Alkanes vs silanes: …