Diamond
Definition and meaning of Diamond in chemistry.
Diamond is a naturally occurring crystalline allotrope of carbon. An allotrope is a different physical form of the exact same element. In diamond, each carbon atom forms four strong covalent sigma bonds to neighboring carbon atoms. This rigid network gives diamond extraordinary physical properties like supreme hardness and high thermal conductivity.
In more detail
Diamond features a three-dimensional network of identical carbon-carbon covalent bonds. These strong bonds extend throughout the entire crystal. This arrangement gives diamond its extreme hardness of 10 on the Mohs scale.
It also creates a high melting point of nearly 4000 degrees Celsius. Diamond has very low electrical conductivity. All its valence electrons are strongly localized in sigma bonds instead of moving freely.
The highly ordered carbon lattice also facilitates rapid heat transport. This gives diamond the highest thermal conductivity of any known natural material. It transfers heat much faster than metallic copper.
Diamond also possesses a high refractive index and significant optical dispersion. These properties result from the dense packing of its atoms and strong bonds. The dense lattice significantly slows down light passing through the crystal.
Natural diamonds form under extreme pressure and high temperature deep in Earth's mantle. Deep-origin volcanic eruptions eventually transport these crystals to the surface. Synthetic diamonds are now produced industrially at a massive scale.
Manufacturers use special high-pressure presses to mimic the extreme mantle conditions. They also use chemical vapor deposition to grow diamond films directly from plasma. A common misconception claims that diamonds formed from compressed coal. In reality, most natural diamonds predate the existence of terrestrial plant life.
Key facts
| Formula | C (carbon) |
|---|---|
| Field | Inorganic Chemistry |
| Hardness | 10 on the Mohs scale |
| Crystal structure | Face-centered cubic |
| Thermal conductivity | Highest of any natural bulk material |
| Band gap | 5.5 eV (excellent electrical insulator) |
Industrial diamond grit consists of the exact same carbon lattice as gem-quality diamonds. Manufacturers bond this hard grit onto circular saw blades and oil well drill bits. Workers use these tough tools to cut and grind extremely hard materials like concrete, granite, and subterranean bedrock. Additionally, engineers use synthetic diamond films in high-power electronics as protective heat spreaders.
Frequently asked questions
Why is diamond so hard while graphite is soft?
Diamond features a rigid, three-dimensional network of strong covalent bonds. Graphite features flat layers of carbon atoms held together by weak forces. These weak forces allow the layers in graphite to slide easily.
Can a diamond burn?
Yes. Diamond is pure carbon and will burn in pure oxygen at high temperatures. It produces carbon dioxide gas and leaves no ash behind.
Are diamonds perfectly stable at room temperature?
No, diamond is technically metastable at standard room temperature and pressure. Graphite is the truly stable phase of carbon. Diamond is very slowly converting into graphite, but the rate is essentially zero.
Related terms
Sources & references
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