Diamond and graphite are two very different forms of carbon.

An element is identified by the number of protons in its atoms. Carbon atoms have six protons whether they occur in a diamond, a graphite crystal or a living cell. That identity does not specify how atoms connect. Different structural forms of the same element are called allotropes, and their properties can differ dramatically.

In diamond, each carbon atom bonds to four neighbors in a three-dimensional network. Strong bonds extend throughout the crystal, helping explain its exceptional resistance to scratching. Hardness is not the same as indestructibility, however: a diamond can chip or fracture because breaking behavior also depends on crystal structure and the direction of applied forces.

In graphite, each carbon bonds strongly to three neighbors within a sheet. Interactions between sheets are much weaker than the bonds inside them, allowing layers to slide and fragments to transfer onto paper. That structure explains why graphite can work as both a writing material and a solid lubricant under appropriate conditions.

The Royal Society of Chemistry lists diamond, graphite, graphene and fullerenes among carbon's allotropes. These are not merely different colors or grades of one mineral. They illustrate how changing the organization of a single element produces solids with different mechanical, optical and electrical behavior, even before other elements are added.

The term "pencil lead" is a historical name, not a statement that the core is metallic lead. Conventional graphite pencils use a mixture, commonly including clay. Changing the recipe changes how hard the core feels and how dark a mark it makes. The carbon comparison applies most precisely to graphite, not the entire manufactured pencil.

Electrical behavior provides another clear contrast. Graphite has electrons that can move relatively readily within its sheets, making it electrically conductive, especially along those layers. Ordinary pure diamond is an electrical insulator because its electrons are held in bonds that do not provide the same mobile charge carriers. Impurities can change that behavior.

Diamond is nevertheless a very good conductor of heat. Heat conduction and electrical conduction need not go together: energy can travel through vibrations of the atomic lattice rather than through freely moving electrons. This distinction is why "conductive" should always be qualified by whether someone means heat or electric current.

Turning graphite into diamond requires rearranging strong bonds, not merely squeezing a pencil by hand. Industrial diamond production can use high pressures and temperatures, while chemical vapor deposition grows diamond from carbon-containing gases under controlled conditions. Neither process implies that a finished pencil is a convenient source of jewelry-quality crystals.

At ordinary pressure, graphite is the more thermodynamically stable form, but diamond does not quickly change into graphite on a desk. A large kinetic barrier slows the rearrangement. Stability describes which state is energetically favored; reaction rate describes how quickly a change happens. Confusing those ideas leads to the mistaken expectation that diamonds should visibly decay.

Chemical identity alone cannot predict a material's performance. Diamond and graphite make an especially clear comparison because their atoms belong to the same element while their bonding arrangements differ. The pencil adds another distinction: a manufactured product can contain an allotrope without being a pure sample of it.