Diamond and graphite are both allotropes of carbon — meaning they are different structural forms of the exact same element (carbon, atomic number 6), differing only in how their carbon atoms are bonded and arranged in space, not in what element they are made of. This is a classic example of how the same element can display very different physical properties purely due to differences in atomic structure/bonding, a concept known as allotropy.

• Diamond: Each carbon atom is sp³ hybridised and forms four strong covalent bonds to four neighbouring carbon atoms, creating a rigid, three-dimensional tetrahedral network that extends throughout the crystal. This continuous strong bonding in all directions makes diamond the hardest known natural substance, gives it a high melting point, makes it transparent to visible light, and makes it a poor conductor of electricity (all valence electrons are locked into strong, localized covalent bonds with no free electrons to carry current).
• Graphite: Each carbon atom is sp² hybridised and forms three covalent bonds within a flat hexagonal layer (like a honeycomb), leaving one unhybridised electron per atom that becomes part of a delocalised π-electron system spread across the layer. The layers themselves are held together only by weak van der Waals forces, allowing them to slide over one another. This gives graphite its softness and slippery, lubricating feel, its opaque grey-black appearance, and — crucially — good electrical conductivity along the layers, because the delocalised electrons are free to move.

Because the atoms in both cases are 100% carbon and no other element is involved, the correct answer is carbon. The other elements listed can be ruled out directly:

• Hydrogen is not part of the diamond or graphite lattice at all; hydrogen forms molecular compounds like water or hydrocarbons, not extended covalent solids of this kind.
• Oxygen similarly is not the building block here — oxygen more commonly forms diatomic O₂ gas or compounds like oxides.
• Nitrogen is also unrelated to these structures; it exists as N₂ gas under normal conditions and is not the atom forming diamond or graphite lattices (though trace nitrogen impurities can occur in natural diamonds, they are not what constitutes the bulk material).

This example is frequently used to teach that identity of an element (same protons, same atomic number) can coexist with drastically different physical properties depending purely on structural arrangement and bonding type — the concept of allotropy.