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Graphite is a much better conductor of h...

Graphite is a much better conductor of heat and electricity than diamond. This is due to the fact that each carbon atom in graphite:( ग्रेफ़ाइट ऊष्मा ओर धारा का चालक होता है , क्यों ?)

A

undergoes `sp^(2)` hybridization and forms three sigma bond with three neighbouring carbon atoms

B

undergoes `sp^(2)` hybridization

C

is tetrahedrally bonded

D

is free from van der Waals force

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The correct Answer is:
**Step-by-Step Solution:** 1. **Understanding the Structure of Graphite and Diamond:** - Graphite and diamond are both allotropes of carbon but have different structures. Graphite has a layered structure, while diamond has a three-dimensional tetrahedral structure. 2. **Hybridization in Graphite:** - In graphite, each carbon atom is sp² hybridized. This means that each carbon atom forms three sigma (σ) bonds with three neighboring carbon atoms, leaving one electron unbonded. 3. **Presence of Delocalized Electrons:** - The unbonded electron from each carbon atom in graphite is free to move. This electron is delocalized over the layers of graphite, allowing it to move easily. 4. **Electrical and Thermal Conductivity:** - The movement of these delocalized electrons is what allows graphite to conduct electricity and heat effectively. The more mobile the electrons, the better the conductivity. 5. **Comparison with Diamond:** - In contrast, in diamond, each carbon atom is sp³ hybridized and forms four sigma bonds with four neighboring carbon atoms. There are no free or delocalized electrons in diamond, which makes it a poor conductor of electricity and heat. 6. **Conclusion:** - Therefore, the reason graphite is a better conductor of heat and electricity than diamond is due to the presence of delocalized electrons resulting from the sp² hybridization of carbon atoms in graphite. **Final Answer:** Graphite is a much better conductor of heat and electricity than diamond because each carbon atom in graphite undergoes sp² hybridization, forming three sigma bonds with neighboring carbon atoms, which allows one electron to be delocalized and move freely. ---
**Step-by-Step Solution:** 1. **Understanding the Structure of Graphite and Diamond:** - Graphite and diamond are both allotropes of carbon but have different structures. Graphite has a layered structure, while diamond has a three-dimensional tetrahedral structure. 2. **Hybridization in Graphite:** - In graphite, each carbon atom is sp² hybridized. This means that each carbon atom forms three sigma (σ) bonds with three neighboring carbon atoms, leaving one electron unbonded. ...
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