What is the reason graphite conducts electricity while diamond does not?

What is the reason graphite conducts electricity while diamond does not?

• A. Graphite has delocalized electrons; diamond lacks them
• B. Graphite has free ions; diamond has fixed ions
• C. Graphite is metallic; diamond is network covalent
• D. Diamond is a liquid

Answer: (A)

Electric conductivity in solids depends on the presence of mobile charge carriers. In graphite, each carbon atom is sp2-hybridized, forming strong bonds with three neighbors in a plane and leaving one electron in a p orbital. These p electrons overlap across the entire layer to form delocalized pi bonds, creating a sea of electrons that can move freely within the sheet. These mobile electrons carry charge when a potential is applied, so graphite conducts electricity along the layers (though conduction between layers is much weaker due to the weaker van der Waals forces between sheets). In diamond, each carbon is sp3-hybridized, making a rigid three-dimensional network of strong covalent bonds. All valence electrons are tied up in these localized bonds, so there are no free charge carriers to move and carry current. That’s why diamond does not conduct electricity. So the key idea is the presence of delocalized electrons in graphite versus the localized electrons in the covalent network of diamond.

Electric conductivity in solids depends on the presence of mobile charge carriers. In graphite, each carbon atom is sp2-hybridized, forming strong bonds with three neighbors in a plane and leaving one electron in a p orbital. These p electrons overlap across the entire layer to form delocalized pi bonds, creating a sea of electrons that can move freely within the sheet. These mobile electrons carry charge when a potential is applied, so graphite conducts electricity along the layers (though conduction between layers is much weaker due to the weaker van der Waals forces between sheets).

In diamond, each carbon is sp3-hybridized, making a rigid three-dimensional network of strong covalent bonds. All valence electrons are tied up in these localized bonds, so there are no free charge carriers to move and carry current. That’s why diamond does not conduct electricity.

So the key idea is the presence of delocalized electrons in graphite versus the localized electrons in the covalent network of diamond.