Diamond and graphite present an interesting paradox: two materials with dramatically contrasting properties—one a clear, brilliant substance prized for its extreme hardness, the other an opaque, dark material used as a soft lubricant—share the exact same chemical basis. This immense difference in physical characteristics, from transparency to density, suggests a fundamental difference in their underlying chemistry. Yet, their composition reveals a surprising identity.
The Core Answer: A Single Element
Both diamond and graphite are composed almost entirely of a single element: Carbon. Carbon is identified on the periodic table by the atomic symbol C and atomic number 6. It is the fourth most abundant element in the universe and forms the backbone of all known organic life on Earth. In its pure, solid form, carbon possesses the unique capacity to arrange its atoms in multiple ways. The primary distinction between the two materials is not a difference in the atoms themselves, but rather in how those atoms are structurally connected.
The Concept of Allotropes
The scientific explanation for this chemical identity and physical difference is the phenomenon of allotropy. An allotrope is one or more forms of a chemical element that can exist in the same physical state. The ability of carbon to form diamond and graphite is a direct result of its atoms bonding together in completely different geometric arrangements. Allotropes maintain the same chemical composition but display unique physical properties because of their distinct internal structures.
Diamond’s Rigid Crystalline Structure and Properties
The specific arrangement of carbon atoms in a diamond is known as a tetrahedral lattice. In this configuration, every carbon atom forms a strong covalent bond with four neighboring carbon atoms. These bonds are oriented in three dimensions, creating an interlocking, continuous network that extends throughout the entire crystal. This dense, uniformly strong framework explains the material’s extreme properties.
The structure results in diamond being the hardest naturally occurring material, registering a 10 on the Mohs scale of hardness. The tight bonding prevents the movement of atoms, which gives it a high density. Furthermore, because all outer electrons are tightly locked into these localized covalent bonds, diamond is an excellent electrical insulator and does not conduct electricity. The interlocking nature of this structure requires a vast amount of energy to break, contributing to its extremely high melting point.
Graphite’s Flexible Layered Structure and Properties
In sharp contrast to diamond, the carbon atoms in graphite are arranged in flat, two-dimensional sheets. Within each sheet, carbon atoms are covalently bonded to only three neighbors, forming a repeating pattern of hexagonal rings. The fourth valence electron is not localized in a bond and is free to move within the layer. These delocalized electrons allow charge to flow easily across the plane of the layer, making graphite an excellent conductor of electricity.
The individual layers, often referred to as graphene sheets, are held together by comparatively weak van der Waals forces. Because these forces are weak, the sheets can easily slide past one another. This sliding motion makes graphite feel slippery and gives it a low Mohs hardness of 1 to 2. This mechanism allows graphite to be used as a lubricant and to leave a mark when a pencil is drawn across paper.