Why can graphite conduct electricity but not diamond?
In graphite, each carbon atom is covalently bonded to three other carbon atoms. One of the four outer shell electrons of each carbon atom is therefore not engaged in bonding, and becomes delocalised. These delocalised electrons are free to move around the structure, carrying charge and allowing graphite to conduct electricity.
However, in diamond, each carbon atom is covalently bonded to four other carbon atoms. All four of the outer shell electrons in each carbon atom are engaged in bonding, leaving no delocalised electrons free to move around the structure and carry charge.
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