Q.Compare the crystal structures of diamond and graphite. In your answer, describe the hybridisation of carbon in each, explain why diamond is extremely hard and an electrical insulator, and why graphite is soft, slippery and electrically conducting.

Graphite is the thermodynamically most stable form of carbon under normal conditions.

Among carbon's allotropes/forms, graphite has the lowest standard Gibbs free energy (and lowest enthalpy) at room temperature and pressure, making it the thermodynamically most stable form — this is why diamond is described as being 'metastable' (kinetically stable but thermodynamically less stable than graphite, converting to graphite only under extreme c …

Fullerene, graphite, and graphene share the same trigonal-planar, sp² carbon bonding of fused six-membered (and, in fullerene, also five-membered) rings; dry ice is solid $CO_2$, whose carbon is sp hybridised in a linear molecule, so it is the odd one out.

Graphite consists of 2-D sheets of sp² carbon atoms, each bonded to three neighbours in a hexagonal honeycomb lattice, with a delocalised pi-electron system above and below each sheet. Graphene is simply a single isolated layer of graphite — chemically identical sp² bonding. Fullerene ($C_{60}$ etc.) is a closed, curved cage built from the same three-sigma-bonds-per-carbon sp² framework, just curved by incorporating some five-membered rings among the hexagons. Dry ice, by contrast, is the solid form of carbon dioxide, $CO_2$, a small linear triatomic molecule ($O=C=O$) in which the central carb …

Diamond is a giant covalent (network) solid, held together entirely by C–C covalent bonds.

In diamond, every carbon atom is sp³ hybridized and forms four strong covalent bonds to four neighbouring carbon atoms in a tetrahedral arrangement, extending throughout the crystal as a three-dimensional network. This is why diamond is extremely hard …

Each carbon in $C_{60}$ fullerene is bonded to exactly three other carbons (forming the closed, ball-shaped network of fused hexagons and pentagons), the same three-connected bonding pattern as planar graphite, so each carbon is $sp^2$ hybridised.

Fullerene $C_{60}$ consists of 60 carbon atoms arranged in a closed, roughly spherical (football/soccer-ball-shaped) cage made of 20 six-membered rings and 12 five-membered rings. Each carbon atom forms exactly three sigma bonds to three neighbouring carbon atoms (as in graphite's 2-D sheets), leaving one unhybridised p-orbital per carbon that contributes to a delocalised pi system over the cage surface. Three …

Graphite's structure is stacked, planar sheets of hexagonally-arranged, sp2-hybridised carbon atoms.

In graphite, every carbon atom is sp2 hybridised and forms three strong sigma covalent bonds to three neighbouring carbon atoms (C-C bond length about 141.5 pm), building a continuous, flat, honeycomb-like network of fused hexagonal rings - essentially an infinite sheet, similar to a single layer of what is now called graphene. Each carbon atom's fourth (unhybridised) electron sits in a p-orbital perpendicular to the sheet; these p-orbitals overlap to give a delocalised pi-electron cloud spread over the whole sheet, which is why graphite conducts electricity along the plane of the layers.

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Graphite is a giant covalent (network) solid — option (b).

From NCERT Class 11 Chemistry (p-Block, carbon allotropes): in graphite each carbon is sp² hybridised and covalently bonded to three neighbours, forming hexagonal sheets; the whole layer is one giant covalent network. The fourth electron is delocalised over the layer (giving electr …

Graphite's electrons are delocalised/spread out across the layers — option (d).

From NCERT Class 11 Chemistry (p-Block, carbon allotropes): in graphite each carbon is sp² hybridised, forming three σ-bonds within a planar hexagonal sheet. The remaining unhybridised p-electron on every carbon forms a delocalised π-system spread over the entire layer. These mobile, delocalised electrons …

Carbon exists in crystalline forms (diamond, graphite) and amorphous forms (lamp black, coke, charcoal) that lack a regular repeating lattice.

Carbon's allotropes are broadly divided into crystalline allotropes, where atoms are arranged in a well-defined repeating three-dimensional lattice (diamond: tetrahedral sp3 network; graphite: layered sp2 sheets), and amorphous allotropes, which consist of very fine crystallites of graphite arranged without long-range order. Lamp bl …

Graphite is the thermodynamically most stable allotrope of carbon at room temperature and pressure.

Carbon exists as diamond, graphite, fullerenes, etc. At ordinary temperature and pressure, graphite has the lowest free energy, so it is the reference (most stable) allotrope; diamond is only metastable and very slowly conver …