CARBON – From the element of life to advanced technology _ PART I (WHY is Carbon special?)

Besides Silicon, another miracle the nature gives us is Carbon. This element is impressive. It is capable of forming many allotropes (structurally different forms of the same element) due to its valency. Well-known forms of carbon include of course the diamond and the graphite, but in recent decades, many more allotropes have been discovered and researched including ball shapes such as buckminsterfullerene and sheets such as graphene. Larger scale structures of carbon also include nanotubes, nanobuds and nanoribbons. Other unusual forms of carbon exist at very high temperatures or extreme pressures. According to the Samara Carbon Allotrope Database there are around 500 hypothetical 3-periodic allotropes of carbon known at the present time.

Few of the most important carbon allotropes

All these carbon allotropes have remarkable properties and graphite is one of the first of them, humans use for a basic but essential skill: handwriting. When we write we do exactly this we transfer atoms from the pencil to the paper. Graphite is an excellent material for artistic expression as writing and drawing. It is important and interesting to note, that although diamond is culturally revered as the superior  form of carbon, it is in fact incapable of deep expression and unlike graphite no good art can come from diamond.

This description of the relationship between the two forms of Carbon, (diamond and graphite), as a rivalry is exactly a battle between dark, expressive, utilitarian graphite and sublime, cool, hard, glinting diamond has been raging since antiquity. In terms of cultural value, diamond has long been the winner, but that may be set to change. A new understanding of graphite’s inner structure has made it a source of wonder.

In 2010 professor Andre Geim one of the world’s foremost carbon experts from Manchester University in the U.K. together with his team had received the 2010 Nobel Prize for Physics for his groundbreaking work on graphene, a 2-dimensional version of graphite and a marvel of the materials world.

Carbon is a light atom with 6 protons and usually 6 neutrons in its nucleus. Sometimes it contains 8 neutrons, but in this form, known as carbon-14, the atomic nucleus is unstable, and so the element falls apart through radioactive decay. Because the rate of this decay is consistent over long period of time, and because this form of carbon finds its way into many materials, measuring its presence in a material allow us to work out that material’s age. This scientific method, known as carbon dating, has thrown more light than any other on our ancient past. The true ages of Stonehenge, the Turin Shroud and the Dead Sea Scrolls have all been revealed by this form of carbon.

Radioactivity aside, the nucleus plays a back-seat role in carbon. In terms of all of its other properties and behavior it is the 6 electrons that surround and shield the nucleus that are important. 2 of these electrons are deeply embedded in an inner core near to the nucleus and play no role in the atom’s chemical life – its interaction with other elements. This leaves 4 electrons, which form its outermost layer, that are active. It is these 4 electrons that make the difference between the graphite of a pencil and the diamond of an engagement ring.


The simplest thing a carbon atom can do is share each of these 4 electrons with another carbon atom, forming 4 chemical bonds. This solves the problem of its active 4 electrons: Each electron is partnered off with a corresponding electron, belonging to another carbon atom. The crystal structure produced extremely rigid. It is a diamond.

Many precious gemstones, including rubies, emeralds, and sapphires, form in the planet’s crust. But, diamonds form hundreds of kilometres below the crust in the Earth’s mantle. The Earth’s mantle is a high temperature, high pressure environment. Carbon atoms that are under these conditions will start to bond together to form crystals, eventually turning into diamonds. It is believed that massive volcanic eruptions are responsible for bringing diamonds close to the Earth’s surface from the depths of the planet’s mantle. The diamonds that make this journey are found within Kimberlite, which is cooled volcanic material. Kimberlite completely preserves the diamond’s natural form so they can be mined from the earth and sold to consumers around the world.

The biggest diamond yet discovered is located in the Milky Way in the constellation of Serpens Cauda, where it is orbiting A pulsar called PSR J1719-1438. It is an entire planet 5 times the size of Earth. Diamonds on Earth are minuscule by comparison. The biggest yet diamond found on Earth is the size of a football. Extracted from the Cullinan mine in South Africa, it was eventually presented to King Edward VII in 1907 on his birthday and is now part of the Crown Jewels of the British monarchy. This diamond was formed far below the surface of the Earth at a depth of approximately 300 km, where, over the course of billions years, the high temperatures and pressures converted a largish-sized carbon rock into the huge diamond. The diamond was then most likely carried to the surface of our planet during a volcanic eruption, where it lay inert and undisturbed for millions of years until it was discovered a mile underground. The entire Cullinan diamond would have weighed little more than half a kilogram.

India was the sole source of diamonds until the mid 18th century, when they were discovered in other parts of the world, most notably South Africa. Each diamond is, in fact, a single crystal. In a typical diamond there are about a million billion billion atoms (1, perfectly arranged and assembled into this pyramidal structure. And it is this structure that accounts for its remarkable properties. In this formation, the electrons are locked into an extremely stable state, and this is what gives it its legendary strength. It is also transparent, but with an unusually high optical dispersion, which means that it splits light that enters it into its constituent colours, giving it its bright rainbow sparkle.The combination of extreme hardness and optical lustre makes diamonds almost flawless as gemstones.

Because of their hardness, virtually nothing can scratch them, and so they keep their perfectly faceted shape and pristine sparkle not just throughout the lifetime of the wearer but throughout the lifetime of a civilization – through rain or shine, whether worn in a sandstorm, hacking through a jungle, or just doing the’ washing up. Even in antiquity diamond was known to be the hardest material in the world. The word ‘diamond’ is derived from the Greek adamas, meaning ‘unalterable’ or ‘unbreakable’.

Unlike gold, diamonds have never been part of the world’s monetary system, despite their financial value. They are not a liquid asset – and quite literally so: they cannot be melted down and, in this way, commodified. Large diamond gemstones have no use except to arouse wonder and awe and, most importantly, to affirm status. Before the 20th century only the truly rich could afford them. But the growing wealth of the European middle classes provided a tempting new market for diamond miners.

The problem faced by the company DeBeers, which in 1902 controlled 90% of the world’s diamond production was how to sell to this much bigger market without devaluing the gems in the process. They managed it through a cunning marketing campaign: by concocting the phrase “diamonds are forever“, they invented the idea of the diamond engagement ring as the only true way to express everlasting love. Anyone who wished to convince their lover of the truth of their feelings needed to buy one, and the more expensive the diamond, the truer the feelings expressed. The marketing campaign took off spectacularly, catapulting a diamond into millions of households and culminating in a James Bond movie, accompanied by a Shirley Bassey/john Barry song, that enshrined the new social role of the diamond as the embodiment of romantic love.

But diamonds are not for ever, at least on the surface of this Planet. It is, in fact, diamond’s sibling structure, graphite, that is the more stable form, and so all diamonds, including the Great Star of Africa in the Tower of London, are actually turning slowly into graphite. This is distressing news for anyone who owns a diamond, although they can be reassured that it will take billions of years before they see an appreciable degradation of their germs.


The structure of graphite is radically different from diamond it consists of planes of carbon atoms connected in a hexagonal pattern. Each plane is an extremely strong and stable structure, and the bonds between the carbon atoms are stronger than those in diamond – which is surprising, given that graphite is so weak that it is used as a lubricant and as lead in pencils.

The conundrum can be explained by noting that within the graphite layers each carbon atom has 3 neighbors with which it shares its 4 electrons. In the diamond structure, each carbon atom shares its 4 electrons with 4 atoms. This gives the individual graphite layers a different electronic structure and stronger chemical bonding than diamond. The flip side, though, is that each atom in graphite has no electrons left over to form strong bonds between its layers. Instead, these layers are held together by the universal glue of the material world, a weak set of forces generated by fluctuations in the electric field of molecules, called van der Waals forces. This is the same force that makes Blu-Tack sticky. The upshot is that when graphite is put under stress, it is the weak van der Waals forces that break first, making graphite very soft.

This is how a pencil works: as you press it on the paper you break the van der Waals bonds and layers of graphite slide across one another, depositing themselves on the page. If it weren’t for the weak van der Waals bonds, graphite would be stronger than diamond.

Take a look at the graphite of a pencil and you will see that it is dark grey and shiny like a metal. For thousands of years it was mistaken for lead and was referred to as “plumbago”, or “black lead”, hence the use of the term “lead” to refer to the graphite used in a pencil. The confusion is understandable since they are both soft metals (although these days we call graphite a semi-metal). Plumbago mines became more and more valuable as new uses were found for graphite, such as the discovery that it was the perfect material to cast cannon and musket balls.

The reason why graphite is metallic, while diamond is not, is also its hexagonal atomic structure. As we have seen, in the diamond structure, all 4 electrons in each carbon atom are partnered up with a corresponding electron. In this way, all atoms in the lattice are strongly held in a bond, and there are no ‘free’ electrons. This is the reason why diamonds do not conduct electricity, because there are no electrons free to move within the structure to carry the electric current.

In the graphite structure, on the other hand, the outer electrons do not just bond with a counterpart electron in a neighboring atom but rather form a sea of electrons within the material. This has several effects, one of which is to allow graphite to conduct electricity, since the electrons can move around like fluid. Graphite was used by Edison for the 1st light bulb filaments because it also has a high melting point, which allows it to glow white hot without melting when a high current passes through it. Meanwhile, the sea of electrons also acts as an electromagnetic trampoline for light, and this reflection of light is what makes it appear shiny like other metals. This neat explanation of graphite’s metallic properties is not what won Andre Geim’s team a Nobel Prize, though. It was merely their starting point.

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