Diamond and graphite: the fascinating world of two allotropes

In our daily life, carbon exists in many forms, the most well-known of which are graphite in pencil leads and dazzling diamonds - diamonds. Although they are derived from the same element, the physical properties of the two are very different, from color, hardness to melting point, showing the diversity and magic of carbon.

Structural differences: understanding macroscopic differences from the microscopic

Diamond and graphite are both made of carbon atoms connected by covalent bonds, but their arrangement is completely different. Diamond is much harder than graphite because the carbon atoms in diamond are arranged in a tetrahedral structure, and each carbon atom is connected to four other carbon atoms, forming an extremely hard and uniform spatial network structure. No matter which direction the external force is applied, a large number of covalent bonds need to be broken at the same time to deform or break it.

 

In contrast, the structure of graphite appears to be much "loose". The carbon atoms in graphite are arranged in layers, and the carbon atoms in each layer are closely connected by covalent bonds to form a hexagonal grid, while the layers are connected to each other by weaker van der Waals forces. The distance between layers is too large and the force is too weak, so it is easy to be "broken one by one" - first it is easily "rubbed" into extremely thin layers, and then the microscopic layer structure is easily destroyed by external forces. This layered structure gives graphite good lubricity and plasticity, making it easy to cut and shape, and its hardness is much lower than diamond.

 

From graphite to diamond: the miracle of artificial synthesis

Given the huge difference between diamond and graphite, scientists have long been committed to exploring methods to synthesize diamond from graphite. From Moissan's high-temperature electric furnace attempt, to the later explosion method, vapor deposition method, and then to the modern high-temperature and high-pressure method, each technological innovation marks the deepening of human understanding of carbon materials and the improvement of technical capabilities. Especially the vapor deposition method and the high-temperature and high-pressure method, the former can grow diamond films or crystals on a specific substrate by precisely controlling the deposition process of carbon atoms; the latter uses the catalytic effect of catalysts under high temperature and high pressure conditions to convert graphite into large particles of diamond, which are used in industrial cutting tools and jewelry.

 

Anomaly of hardness and melting point: Why does diamond have a low melting point?

From a microscopic perspective, melting means that the particles that make up the substance gain freedom in three-dimensional space and can flow freely. For diamond and graphite, this freedom requires the simultaneous destruction of a large number of covalent bonds, so their melting points are very high.

 

For most crystals, the higher the hardness, the higher the melting point. However, in the case of diamond and graphite, the hardness and melting point are inconsistent.

 

Although diamond is known for its unparalleled hardness, its melting point is unexpectedly lower than that of graphite. The reason behind this is closely related to their covalent bond strength and structural characteristics. The carbon atoms in diamond use sp3 hybridization, and the covalent bond length formed is longer (0.155nm) and the bond energy is relatively low; while the carbon atoms in graphite use sp2 hybridization, the bond length is shorter (0.142nm) and the bond energy is higher. Therefore, when both materials transform from solid to liquid, although a large number of covalent bonds need to be broken, the stronger covalent bonds in graphite require higher energy to break, resulting in a higher melting point for graphite than for diamond (3680°C for graphite and 3550°C for diamond).

 

Thermal Conductivity of Graphite and Diamond

Graphite is a material with excellent thermal conductivity, and its thermal conductivity is much higher than many common materials. The thermal conductivity range of graphite is generally high, but the specific value varies depending on the quality of the graphite and the test conditions.

 

The layered structure of graphite is the key to its efficient thermal conductivity. The carbon atoms in the layers are tightly bound by strong covalent bonds to form a stable structure, which is conducive to the rapid transfer of heat. However, because the layers are connected by weak van der Waals forces, the thermal conductivity of graphite in the interlayer direction is relatively weak. Despite this, graphite is still widely used as a thermal management material in high-temperature environments, such as heat sinks, thermal conductive films, etc. Its excellent thermal conductivity and chemical stability play an important role in these applications.

 

For diamond, although diamond is an insulator and does not contain free electrons, it has the best thermal conductivity of all solids. Its thermal conductivity ranks among the best in nature. ‌At room temperature, the thermal conductivity of diamond can reach 2000~2200 W/(m·K), which is 4~5 times that of copper and silver, 4 times that of silicon carbide (SiC), 13 times that of silicon (Si), and 43 times that of gallium arsenide (GaAs). In addition, the thermal conductivity of type IIa diamond at liquid nitrogen temperature can reach 25 times that of copper, showing super thermal conductivity. Diamond has stable chemical properties, is resistant to acids and alkalis, and does not react with certain chemicals at high temperatures. These properties enable it to maintain good thermal conductivity even in extreme environments.

 

There are no free electrons in the diamond structure, so how can it have thermal conductivity? It turns out that the essence of thermal conductivity and electrical conductivity is different, which is determined by the microscopic nature of heat - the microscopic essence of heat is the movement of particles. If the movement rate of microscopic particles is fast, the external manifestation is high temperature. This movement of microscopic particles can be free and irregular, or it can be self-vibration on the lattice. It can be imagined that the excellent thermal conductivity of diamond is achieved by the vibration of the carbon atoms themselves on the lattice. Due to the highly ordered arrangement of the diamond lattice, and the fact that its vibration frequency is highly consistent with the frequency required for heat (essentially an electromagnetic wave) conduction, this vibration of carbon atoms can easily cause resonance in the crystal, thereby quickly conducting heat from one place to another, making diamond the solid substance with the best thermal conductivity.

 

This unique thermal conductivity makes diamond widely used in high-tech fields. For example, in semiconductor chip packaging, diamond can quickly conduct heat to prevent the chip from performing poorly or reducing reliability due to excessive temperature. In addition, diamond is also used to manufacture heat sinks and high thermal conductivity interface materials for high-power electronic devices. Due to its high thermal conductivity and low thermal expansion coefficient, it can effectively reduce the dimensional change of the material when the temperature changes, and improve the stability and reliability of the equipment.

As allotropes of carbon, diamond and graphite show completely different macroscopic properties through their unique microstructures. From their mutual transformation to anomalous physical properties, each discovery is a profound revelation of the mysteries of nature and a testimony to human wisdom and technological progress.