Why is graphite more easily formed than diamond?

Answer

Dear Hannes,

Let’s start with a little more context:

context
Despite the fact that diamond and graphite consist of the same building blocks, namely carbon atoms, both substances differ greatly in their properties. For example, graphite is black and easy to scratch, while diamond is transparent and extremely hard. You should try scratching a diamond with something that isn’t a diamond itself! The reason for this lies in the way in which these building blocks are arranged among themselves. I added some figures to my answer to illustrate – just add them.

Diamond has a 3D structure where each carbon atom is connected to four other carbon atoms. In the figure you see some carbon atoms that seem to have only two (or one) bonds, but that’s not the case: I’m only showing you a small part of the whole. The entire structure of diamond is gigantic; it repeats continuously in all three dimensions.
Graphite looks completely different. It has a layer texture that is a bit difficult to draw. In the figure at the top you see such a layer viewed from above. Also this time I show you only a small part of the whole. We call such a layer graphene, place several on top of each other and we call it graphite. At the bottom of the figure you can see how those layers position themselves in relation to each other when you look from the side.

First question: 3 or 4 bindings?
And now we come to the gist of your first sub-question: you could say that a carbon atom must have four bonds because it has four unpaired electrons. However, in the graphite diagram you only see three bonds – how is this possible?

The answer is that this diagram is a simplification of reality: it shows the arrangement of the atoms, but does not properly visualize all bonds. In graphite, each carbon atom uses three of its electrons to form covalent bonds with its three closest neighbors. This leaves a fourth free electron per carbon atom. However, that electron does not just float there, but does something special that we call delocalization: the free electrons can no longer be associated with a particular atom or pair of atoms and move freely throughout the entire layer. Delocalization also occurs in a molecule such as benzene or in metal bonds.
This means that in addition to those strong covalent bonds, the carbon atoms in such a graphite layer are held together by this ‘cloud’ of free electrons. You can think of the cloud as an extra bond – a bond that ensures that the carbon atoms in such a layer of graphite are bonded even more strongly than the carbon atoms in diamond!

The softness of graphite is caused by the fact that the layers are held together by Vanderwaals forces. Those forces are many times weaker than the covalent bonds discussed above. This is why when you write with a pencil, you actually rub off layer by layer of graphite on your paper.

Second question: energy!
We discussed above the very different structure of graphite and diamond, and how strong the bonds are between the different carbon atoms. Although it sounds very interesting to convert graphite to diamond, it is not nearly that easy! To change from the graphite structure to the diamond structure you have to break the very strong bonds between the atoms and make sure that the atoms start to orient themselves in a different way. So a lot of energy is needed for that. Converting graphite to diamond is possible at very high pressures, when the carbon is dissolved in liquid rock or metal, and when the conditions are right to form diamond. These conditions can be found in a so-called phase diagram, as you can find a simple version of it here. Take it for a while.

Starting from the bubble at the bottom left, you can see in the diagram that by increasing the pressure (up in the graph), we can convert graphite into diamond. By increasing the temperature (to the right in the graph), we quickly arrive at gaseous or liquid carbon. The graph doesn’t show it very clearly, but you can assume that at room temperature we need more than 150,000 times the atmospheric pressure to succeed in our setup.

There is a lot more to say about this, but I think you’ve already had enough of this. Hopefully I could bring some clarity.

Sincerely,
Giele Van den Berghe