Carbon is black (in most cases), non-toxic (a lot of us eat some for breakfast with butter) and doesn’t easily react with other chemicals. An unassuming element. In fact its name comes from the Latin ‘carbo’ meaning ‘coal’. Literally a dull atom.
So now you are probably asking yourself, “Why the f—k are you writing an article about it then?”.
A good question. Short answer: It seemed like a good idea at the time.
It is the key to all life, after all. You and I would not exist without Nature’s humble black ‘Lego brick’.
So let’s give it a go and see what we dig up. You never know, it might be interesting.
The Carbon Atom And Its Two Natural Forms
Fig 1 shows the carbon atom’s ‘components’. The nucleus (the centre of an atom) contains six protons (blue), six neutrons (red). There are six electrons (green) ‘orbiting’ the nucleus. These electrons orbit in two shells. Two are on the inner shell and four on the outer.
The outer four electrons are called the valence electrons (as with all atoms). They can ‘glue together’ a carbon atom with other atoms by ‘sharing’. This sharing of electrons with other atoms is called covalent bonding.
Carbon is something called an allotrope. It’ll ‘stick’ to itself in various ways forming substances with different physical and chemical properties. In other words it can ‘make different stuff’ out of itself.
OK, so I’ve probably lost you at this point and you are racing ahead to the comments (that you should not be reading). I wouldn’t blame you.
For those still with me let’s plough on. Carbon forms two natural allotrope: Diamond and Graphite. (Fig 2)
Diamond and graphite, even though they’re only made from the carbon atom, certainly look different.
In Fig 2 the ‘wire-frame’ drawings are the molecular structure of the two respectively.
The little ‘black dots’ are carbon atoms. The solid black lines joining them together are covalent bonds.
The vertical ‘dashed’ lines in graphite are weak bonds known as Van der Waals bonds. It’s why graphite ‘cleaves’ into sheets rather like slate (basal cleavage).
Diamond has only covalent bonds making it strong in all directions. It doesn’t mean that it can’t be ‘cut’. Diamond cuts in four planes (octahedral cleavage). It’s this fact that a ‘diamond cutter’ uses to create jewellry. (Miller indices.)
Both are made from a hexagonal arrangement of carbon atoms. In the case of graphite they’re arranged 2-dimensionally, flat (dihexagonal) .
For diamond they’re arranged 3-dimensionally (hexoctahedral). Or ‘diamond-cubic’. (Hat-tip to OT.)
This atomic structure makes diamond very hard. Therefore, it is used in drilling and cutting tools.
Diamond also has a high ‘index of refraction’. It more than halves the speed of light, which makes it useful in specialist optics. This is what gives diamonds their sparkle.
Before we move on to carbon based molecules I’d like to tarry-a-while on graphite.
Unlike diamond, graphite conducts electricity. The chances are it’s helping to power the electronics you are using right now. The negative electrode in your mobile device’s lithium ion battery is usually made from graphite.
If we take just one ‘layer’ of graphite (Fig 2 has three) we get a substance called Graphene. The ‘stuff’ of nano-technology!
It’s transparent, strong (200 times that of steel), conducts electricity and is strongly diamagnetic. This means that it creates a magnetic field opposite to the one applied.
If you were to place a ‘sheet’ of Graphene above a powerful magnet it would float.
Unfortunately Graphene is hard to make. Although people have been making small amounts of it for centuries – thanks to the humble pencil. Just by drawing with a pencil the chances are that you are making a tiny amount of Graphene!
If you wish to make your own ‘Back To The Future’ hover-board you can buy something called ‘Pyrolytic graphite’ from eBay (last time I checked). Just add a powerful neodymium magnet and it’ll float in mid-air forever!
(In Part 2 we deal with some simple molecules that carbon forms with other atoms.)
(C) Dr Mike Finnley