**2. Diamond basic**

Diamond, with a Vickers hardness of more than 100 GPa, is one of the most commonly used superhard materials for industrial applications because it has short and strong covalent bonds between carbon atoms in cubic crystal structures. It is a metastable allotrope of carbon, with atoms arranged in a modified surface-centered cube (fcc) crystal structure called a "diamond cube" or "diamond lattice." Currently, there are two approaches, either using graphite as a starting powder, or using diamond in the industry to produce catalyst-free or binderless PCDs, and two carbon allotropes – diamond and graphite (**Figure 3**) will be focused on, which have very different mechanical properties from superhard and supersoft materials, respectively.

**Figure 3.**

*Two carbon allotropes - diamond (left) and graphite (right) [9].*

Diamonds are the hardest known phases in natural and synthetic materials, with a face-centered cubic lattice (space group #227, a = 3.56679 Å). In this compact structure (density of 3.51 g/cm3), the atoms are covalently bonded to four other atoms due to sp3 orbital hybridization, with an average bond length of about 1.54 Å. Graphite has a hexagonal structure (space group #194, a = b = 2.456 Å, c = 6.696 Å), formed by stacks of layers composed of regular hexagonal carbon. Each atom of a given layer is covalently bonded to three other atoms to form a sp2 hybridization type, while the bonds between the layers are weaker, by van der Waals type. The stable form of crystalline carbon at standard temperature and pressure (STP) is graphite, while at higher temperatures and pressures above the well-known Berman-Simon line [10], the cubic form of diamond is stable. Because the diamond is processed using the high temperature and pressure technology above the Berman-Simon line, the diamond retains its unique superhard cube structure when lowered to STP. The formation of diamond from graphite is only a phase transition under HPHT, given by the following equation:

$$\mathbf{C}\_{\text{graphite}} \prec = \mathbf{C}\_{\text{dilamond}} \tag{1}$$

At atmospheric pressure, the Gibbs energy change from graphite to diamond at all temperatures is greater than zero, which means that graphite is a stable phase at all temperatures at atmospheric pressure. However, since diamonds are a denser form of carbon, one would expect increased pressure to make diamond formation more likely. Both diamond powder and graphite powder can be used as raw materials to create bulky, superhard-cutting structures for drilling tools. However, the sintering mechanism is completely different, which has a great impact on the properties of the material.

Due to its extremely high hardness, diamond is widely used in industrial applications that require very high wear resistance. It can be processed as a single crystal in the form of large blocks, or as a powder for grinding, polishing, and grinding operations, or as a PDC tool, where individual grains are joined together by a high pressure and high temperature (HPHT) process. Until the 1950s, only natural diamonds could be used for niche abrasive applications. With the introduction of the HPHT process for synthetic PCDs by GE in 1955 [11], the use of synthetic diamond abrasive materials has increased dramatically. While the performance of natural diamonds varies greatly due to differences in defect density and impurity
