Diamond is a mineral form of the element carbon with its atoms arranged in a crystal structure called diamond cubic. Diamond is a tasteless, odorless, strong, brittle solid, a poor conductor of electricity, colorless in pure form, and insoluble in water. Another solid form of carbon known as graphite is the chemically stable form of carbon at room temperature and pressure, but diamond is metastable and converts to it at a negligible rate under those conditions. Diamond has the highest hardness and thermal conductivity of any natural material, properties that are used in major industrial applications such as cutting and polishing tools.
Because the arrangement of atoms in diamond is extremely rigid, few types of impurity can contaminate it (two exceptions are boron and nitrogen). Small numbers of defects or impurities—about one per million of lattice atoms—can color a diamond blue (boron), yellow (nitrogen), brown (defects), green (radiation exposure), purple, pink, orange, or red. Diamond also has a very high refractive index and a relatively high optical dispersion.
Most natural diamonds have ages between 1 billion and 3.5 billion years. Most were formed at depths between 150 and 250 kilometres (93 and 155 mi) in the Earth's mantle, although a few have come from as deep as 800 kilometres (500 mi). Under high pressure and temperature, carbon-containing fluids dissolved various minerals and replaced them with diamonds. Much more recently (hundreds to tens of million years ago), they were carried to the surface in volcanic eruptions and deposited in igneous rocks known as kimberlites and lamproites.
Synthetic diamonds can be grown from high-purity carbon under high pressures and temperatures or from hydrocarbon gases by chemical vapor deposition (CVD). Natural and synthetic diamonds are most commonly distinguished using optical techniques or thermal conductivity measurements.
Etymology, earliest use and composition discovery
The name diamond is derived from Ancient Greek: ἀδάμας (adámas), 'proper, unalterable, unbreakable, untamed', from ἀ- (a-), 'not' + Ancient Greek: δαμάω (damáō), 'to overpower, tame'. Diamonds are thought to have been first recognized and mined in India, where significant alluvial deposits of the stone could be found many centuries ago along the rivers Penner, Krishna, and Godavari. Diamonds have been known in India for at least 3,000 years but most likely 6,000 years.
Diamonds have been treasured as gemstones since their use as religious icons in ancient India. Their usage in engraving tools also dates to early human history. The popularity of diamonds has risen since the 19th century due to increased supply, improved cutting and polishing techniques, growth in the world economy, and innovative and successful advertising campaigns.
In 1772, the French scientist Antoine Lavoisier used a lens to concentrate the rays of the sun on a diamond in an atmosphere of oxygen, and showed that the only product of the combustion was carbon dioxide, proving that diamond is composed of carbon. Later, in 1797, the English chemist Smithson Tennant repeated and expanded that experiment. By demonstrating that burning diamond and graphite releases the same amount of gas, he established the chemical equivalence of these substances.
Diamond is a solid form of pure carbon with its atoms arranged in a crystal. Solid carbon comes in different forms known as allotropes depending on the type of chemical bond. The two most common allotropes of pure carbon are diamond and graphite. In graphite, the bonds are sp2 orbital hybrids and the atoms form in planes, with each bound to three nearest neighbors, 120 degrees apart. In diamond, they are sp3 and the atoms form tetrahedral, with each bound to four nearest neighbors. Tetrahedra are rigid, the bonds are strong, and, of all known substances, diamond has the greatest number of atoms per unit volume, which is why it is both the hardest and the least compressible. It also has a high density, ranging from 3150 to 3530 kilograms per cubic metre (over three times the density of water) in natural diamonds and 3520 kg/m3 in pure diamond. In graphite, the bonds between nearest neighbors are even stronger, but the bonds between parallel adjacent planes are weak, so the planes easily slip past each other. Thus, graphite is much softer than diamond. However, the stronger bonds make graphite less flammable.
Diamonds have been adopted for many uses because of the material's exceptional physical characteristics. It has the highest thermal conductivity and the highest sound velocity. It has low adhesion and friction, and its coefficient of thermal expansion is extremely low. Its optical transparency extends from the far infrared to the deep ultraviolet, and it has high optical dispersion. It also has high electrical resistance. It is chemically inert, not reacting with most corrosive substances, and has excellent biological compatibility.
The equilibrium pressure and temperature conditions for a transition between graphite and diamond are well established theoretically and experimentally. The equilibrium pressure varies linearly with temperature, between 1.7 GPa at 0 K and 12 GPa at 5000 K (the diamond/graphite/liquid triple point). However, the phases have a wide region about this line where they can coexist. At standard temperature and pressure, 20 °C (293 K) and 1 standard atmosphere (0.10 MPa), the stable phase of carbon is graphite, but diamond is metastable, with a significant kinetic energy barrier that the atoms must overcome to reach the lower energy state, and its rate of conversion to graphite is negligible, with a timescale of millions to billions of years. However, at temperatures above about 4500 K, diamond rapidly converts to graphite. Experiments have found that diamond, in the presence of H2O passes through an intermediate linear carbon phase.
Rapid conversion of graphite to diamond requires pressures well above the equilibrium line: at 2000 K, a pressure of 35 GPa (about 350,000 standard atmospheres) is needed.
Above the graphite–diamond–liquid carbon triple point, the melting point of diamond increases slowly with increasing pressure, but at pressures of hundreds of GPa, it decreases. At high pressures, silicon and germanium have a BC8 body-centered cubic crystal structure, and a similar structure is predicted for carbon at high pressures. At 0 K, the transition is predicted to occur at 1100 GPa.
Results published in Nature Physics in 2010 suggest that, at ultra-high pressures and temperatures (about 10 million atmospheres or 1 TPa and 50,000 °C), diamond melts into a metallic fluid. The extreme conditions required for this to occur are present in the ice giant planets Neptune and Uranus, both of which are made up of approximately 10 percent carbon and could hypothetically contain oceans of liquid carbon. Since large quantities of metallic fluid can affect the magnetic field, this could serve to explain why the geographic and magnetic poles of the two planets are not aligned.
The most common crystal structure of diamond is called diamond cubic. It is formed of unit cells stacked together. Although there are 18 atoms in the figure, each corner atom is shared by eight unit cells and each atom in the center of a face is shared by two, so there are a total of eight atoms per unit cell. The length of each side of the unit cell is denoted by a and is 3.567 angstroms.
The nearest neighbor distance in the diamond lattice is 1.732a/4 where a is the lattice constant, usually given in Angstrøms as a = 3.567 Å, which is 0.3567 nm.
A diamond cubic lattice can be thought of as two interpenetrating face-centered cubic lattices with one displaced by 1⁄4 of the diagonal along a cubic cell, or as one lattice with two atoms associated with each lattice point. Viewed from a <1 1 1> crystallographic direction, it is formed of layers stacked in a repeating ABCABC ... pattern. Diamonds can also form an ABAB ... structure, which is known as hexagonal diamond or lonsdaleite, but this is far less common and is formed under different conditions from cubic carbon.