In this blog post, we’ll examine—from the perspective of materials engineering—why graphite, used in pencil leads, and diamonds, a classic gemstone, exhibit completely different properties even though they’re both composed of the same element: carbon.
Graphite, found in pencils—which almost everyone, not just students, has used at least once—and diamond, well-known as a hard and beautiful gemstone, appear to have absolutely nothing in common at first glance. Graphite is black, soft, and leaves marks easily, whereas diamond can form transparent crystals and possesses extremely high hardness. The difference between the two materials is also quite significant in terms of price and application. So why are we discussing these two such different materials together? Surprisingly, both graphite and diamond are substances composed of carbon atoms. Although they share the same constituent element, their physical properties differ greatly. Understanding this difference helps us grasp an important principle of materials engineering: the properties of a material are not determined solely by the elements it contains, but also vary significantly depending on how the atoms are arranged and bonded to one another.
Graphite is an important carbon material widely used in everyday life and industry. In pencil leads, graphite is mixed with clay and other substances; due to its layered structure, fine particles are released when rubbed against paper, leaving a mark. Additionally, graphite’s ability to conduct electricity and heat makes it useful in various industrial applications, including electrodes, refractory materials, and lubricants. On the other hand, diamonds are not only highly valuable as gemstones but are also utilized in various industrial fields—such as cutting and grinding tools, precision machining, and thermal management materials—due to their extremely high hardness and excellent thermal conductivity. As shown, although graphite and diamond differ greatly in appearance, physical properties, and applications, they are fundamentally composed of carbon atoms. So why do these two substances, both made of carbon, exhibit such completely different characteristics?
In materials engineering, phases and phase transitions are crucial concepts for understanding the properties of materials. A phase transition refers to the phenomenon in which a material’s structure and state change into a different phase as conditions such as temperature, pressure, and composition vary. A familiar example is water. Ice, liquid water, and water vapor all have the same chemical formula, H₂O, but exist in different states depending on temperature and pressure conditions. We can understand the relationship between graphite and diamond from a similar perspective. Although both substances are allotropes composed of the same carbon element, the arrangement of carbon atoms and their bonding patterns differ. However, it is not accurate to simply distinguish between them by saying that graphite is “formed by temperature” and diamond is “formed by pressure.” The structure in which carbon exists stably is influenced by various conditions, including temperature and pressure. Under normal pressure, graphite is the thermodynamically stable form of carbon, while diamond becomes the stable structure under high-pressure conditions. Graphite can also be produced artificially through a graphitization process, in which carbon is treated at very high temperatures. Meanwhile, in nature, diamonds form under the high pressure and temperature conditions found deep within the Earth; artificially, they can be grown using the High-Pressure, High-Temperature (HPHT) method at pressures of approximately 5–6 GPa and temperatures of approximately 1,300–1,600°C. Currently, Chemical Vapor Deposition (CVD), which involves separating carbon atoms from a carbon-containing gas and growing them on a diamond seed crystal, is also widely used. The CVD method holds significant technical importance because it allows diamonds to be grown at much lower pressures than the HPHT method. Ultimately, graphite and diamond are prime examples illustrating that the properties of a material are determined not simply by its constituent elements, but by the conditions under which carbon atoms form a particular crystal structure.
The difference in properties between graphite and diamond can be more clearly understood by examining their atomic bonding structures. In diamond, each carbon atom forms strong covalent bonds with four neighboring carbon atoms. The carbon atoms are arranged in a tetrahedral configuration with a bond angle of approximately 109.5 degrees, and this structure extends continuously throughout three-dimensional space. In other words, within a diamond, carbon atoms form a dense three-dimensional lattice through strong covalent bonds. Because of this strong bonding structure, diamond exhibits extremely high hardness. In contrast, in graphite, a single carbon atom bonds with three neighboring carbon atoms to form a planar structure consisting of interconnected hexagons. Each carbon atom forms an sp² hybrid bond, and the remaining electrons do not remain bound exclusively between specific pairs of atoms but instead form a delocalized π electron system within the layer. These electrons are one of the key reasons graphite exhibits electrical conductivity. Although the carbon atoms within each layer are connected by strong covalent bonds, the forces between layers are relatively weak. Consequently, the layers can easily slide past one another or peel away, which is why graphite particles remain on paper when a pencil is rubbed against it. In the past, it was sometimes simplistically explained that the remaining electrons directly connect the layers of the flaky structure; however, it is more accurate to distinguish between the delocalized electrons within each layer and the weak interlayer interactions. As this example shows, even materials composed of the same carbon atoms can exhibit significantly different physical properties—such as hardness, electrical properties, and optical characteristics—depending on the arrangement of the atoms and the nature of their bonds.
Thus, controlling a material’s structure and manufacturing conditions to produce materials with desired properties is one of the core research areas in materials engineering. Graphite and diamond are prime examples that illustrate this relationship well. This raises an interesting question: If diamonds can be produced artificially, would they still be considered a rare and expensive material? In the past, it was extremely difficult to achieve the high pressures and temperatures required to create synthetic diamonds and to precisely control their quality; there were also significant technical limitations to growing diamonds in the sizes and quality suitable for use as gemstones on a large scale. However, the situation has changed significantly today. With advances in HPHT and CVD technologies, high-quality lab-grown diamonds suitable for jewelry are now being commercially mass-produced, and their size, color, and clarity have also steadily improved. In the CVD process, in particular, a gas containing carbon is introduced into a vacuum chamber and activated so that carbon atoms grow into crystals on a diamond seed crystal. Consequently, the argument that “diamond prices remain high because lab-grown diamonds cannot be mass-produced,” as was commonly held in the past, no longer aligns with the current reality. While both natural and lab-grown diamonds are composed of crystalline carbon and share very similar fundamental chemical and physical properties, they differ in their formation processes, supply structures, scarcity, and market valuation. Today, the challenges in materials engineering extend beyond simply producing diamonds; they now focus on growing larger, more uniform, and defect-free diamonds at lower costs and with less energy consumption, while precisely tailoring their properties to meet the needs of various fields such as electronics, optics, thermal management, and quantum technology.
If materials engineering advances to the point where we can precisely control the arrangement and bonding states of atoms, we will be able to design materials with different properties more freely, even while using the same elements. Of course, this cannot be viewed as entirely identical to the alchemy of the past. Traditional alchemy included attempts to transform one element into another, but transforming graphite into diamond does not involve changing the element carbon itself into a different element; rather, it involves altering the arrangement and bonding structure of carbon atoms. Nevertheless, the fact that both ordinary, black graphite and dazzling, hard diamond are composed of carbon clearly illustrates the fascinating world that materials engineering reveals. Just as important as what elements a material is composed of is how those atoms are arranged and bonded; a key role of materials engineering is to create new materials with desired properties by understanding and controlling precisely that structure. Hidden between the pencil lead we encounter in everyday life and the dazzling gemstone known as the diamond lies the fascinating field of materials engineering—encompassing atomic-level structure and bonding, phase stability, and materials manufacturing technology.
