Structure of Graphite
Graphite has a giant covalent structure in which: each carbon atom is joined to three other carbon atoms by covalent bonds. the carbon atoms form layers with a hexagonal arrangement of atoms. the layers have weak forces between them.
Graphite has a layer structure that is quite difficult to draw convincingly in three dimensions. The diagram below shows the arrangement of the atoms in each layer and the way the layers are spaced.
Solid carbon comes in different forms known as allotropes depending on the type of chemical bond. The two most common are diamond and graphite.
In diamond the bonds are sp3 and the atoms form tetrahedra with each bound to four nearest neighbors. In graphite, they are sp2 orbital hybrids and the atoms form in planes with each bound to three nearest neighbors 120 degrees apart.
The individual layers are called graphene. In each layer, the carbon atoms are arranged in a honeycomb lattice with a bond length of 0.142 nm, and the distance between planes is 0.335 nm. Atoms in the plane are bonded covalently, with only three of the four potential bonding sites satisfied.
The fourth electron is free to migrate in the plane, making graphite electrically conductive. Bonding between layers is via weak van der Waals bonds, which allow layers of graphite to be easily separated, or to slide past each other. Electrical conductivity perpendicular to the layers is consequently about 1000 times lower
Notice that you can’t really draw the side view of the layers to the same scale as the atoms in the layer without one or another part of the diagram is either very spread out or very squashed.
In that case, it is important to give some idea of the distances involved. The distance between the layers is about 2.5 times the distance between the atoms within each layer.
The layers, of course, extend over huge numbers of atoms – not just the few shown above.
You might argue that carbon has to form 4 bonds because of its 4 unpaired electrons, whereas in this diagram it only seems to be forming 3 bonds to the neighboring carbons. This diagram is something of a simplification and shows the arrangement of atoms rather than the bonding.
The Bonding in Graphite
Each carbon atom uses three of its electrons to form simple bonds to its three close neighbors. That leaves the fourth electron at the bonding level. These “spare” electrons in each carbon atom become delocalized over the whole of the sheet of atoms in one layer.
They are no longer associated directly with any particular atom or pair of atoms but are free to wander throughout the whole sheet. The important thing is that the delocalized electrons are free to move anywhere within the sheet – each electron is no longer fixed to a particular carbon atom.
There is, however, no direct contact between the delocalized electrons in one sheet and those in the neighboring sheets. The atoms within a sheet are held together by strong covalent bonds – stronger, in fact, than in diamond because of the additional bonding caused by the delocalized electrons.
So, what holds the sheets together? In graphite, you have the ultimate example of van der Waals dispersion forces. As the delocalized electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below – and so on throughout the whole graphite crystal.
Graphite has many properties like a high melting point, similar to that of a diamond. In order to melt graphite, it isn’t enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure.
It has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. You can think of graphite rather like a pack of cards – each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper.
Graphite has a lower density than diamond. This is because of the relatively large amount of space that is “wasted” between the sheets.
Graphite is insoluble in water and organic solvents – for the same reason that diamond is insoluble. Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite.
It conducts electricity. The delocalized electrons are free to move throughout the sheets. If a piece of graphite is connected to a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.
