Science (086) — AI-generated practice question

Covalent compounds are generally poor conductors because electron sharing produces no free charged particles (ions). However, graphite is an exception due to its unique structure.

In graphite, each carbon atom is bonded to three other carbon atoms in the same plane in a hexagonal arrangement, with one double bond. This leaves each carbon atom with one electron not used in bonding. These electrons are free to move across the layers, allowing graphite to conduct electricity.

Thus, graphite conducts electricity not because it forms ions, but because of its special layered structure that allows the movement of free electrons.

Source: Chapter 4, Allotropes of Carbon (More to Know)

---

• The examiner wants you to first acknowledge the general rule (covalent compounds don't conduct), then explain graphite as an exception using its structure.
• Key facts to mention: 3 bonds per carbon atom → one free/delocalized electron per atom → free electrons carry charge.
• The passage states graphite is "a very good conductor of electricity unlike other non-metals" — always quote or reference the textbook reason.
• Do not use the word "delocalized" if you haven't been taught it formally; instead say "free electrons" which is sufficient for CBSE Class 10.