Diamond is a bad conductor of electricity because its bonding electrons are used in

Step 1: Analyze the question We are asked why diamond does not conduct electricity. The key is to ask what its bonding electrons are doing.

Step 2: Explain the concept Electrical conductivity needs mobile charge carriers, meaning free electrons or free ions.

• Metals conduct because they have delocalised electrons that move freely (metallic bonding).
• Graphite conducts because each carbon uses only 3 of its 4 outer electrons in bonding, leaving one delocalised electron per atom.
• Diamond is different. Each carbon atom uses all 4 of its outer electrons to form 4 strong covalent bonds with 4 other carbon atoms in a rigid tetrahedral arrangement.

Step 3: Apply to diamond Because all the valence electrons are held in localised covalent bonds, there are no free electrons and no ions. No charge carriers can move, so diamond is a poor conductor (an insulator).

Step 4: Check the options

• A. Crystal lattice formation: diamond does have a giant lattice, but a lattice is the structure, not the type of bonding that uses up the electrons. Metals also have lattices and they conduct.
• B. Covalent bond formation: correct. The electrons are all tied up in covalent bonds.
• C. Metallic bond formation: wrong, since diamond has no metallic bonding.
• D. Coordinate bond formation: wrong, since diamond's bonds are normal shared-pair covalent bonds, not dative bonds.

Answer: B. Covalent bond formation

Exam tip: Remember "Diamond: 4 bonds, 0 free electrons → no conduction. Graphite: 3 bonds, 1 free electron → conducts." Questions comparing diamond and graphite appear often in JAMB.