Diamond and graphite are allotropes of carbon, meaning they are made solely of carbon atoms but arranged differently. This structural difference leads to vastly different properties. Diamond has a tetrahedral structure where each carbon is bonded to four others, with all valence electrons tightly held in covalent bonds, making it very hard and a poor conductor of electricity. Graphite has a layered structure where each carbon is bonded to three others, forming hexagonal rings within layers. One valence electron per carbon atom is delocalized and free to move across the layers, making graphite soft and a good conductor of electricity. Therefore, the statement that diamond is a better electric conductor than graphite is false. Options A and B are true: diamond is much harder than graphite, and both are pure carbon, hence containing the same percentage of carbon (100%). Option D is generally considered true in that the distances between non-bonded carbon atoms vary depending on their relative positions in the lattice, unlike the consistent bond length. However, the distinct electrical conductivity is the most prominent difference.