Chemguide: Core Chemistry 14 - 16

Giant covalent structures

This page looks at the way some atoms arrange themselves into giant covalent structures, and the effect this has on their simple properties. It includes silicon dioxide (silica), and carbon as diamond, graphite and graphene.

I am assuming that you already know about covalent bonding.

What is a giant covalent structure?

A giant covalent structure is one in which the atoms are joined up by covalent bonds over huge (but variable) numbers of atoms. It is not a molecule, because the number of atoms joined up in a diamond, say, is completely variable - depending on the size of the crystal.

You may find that these structures are sometimes described as "giant molecular structures". I think this is misleading. A molecule contains a fixed number of atoms (even if that number is very large). In a giant covalent structure like diamond the number of carbon atoms is entirely variable.

The giant covalent structure of diamond

Carbon has an electronic arrangement of 2,4. In diamond, each carbon atom shares electrons with four other carbon atoms - forming four single bonds.

In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but that's because we are only showing a small part of the whole structure. Every carbon atom, until you get a crystal edge, will be joined to four others.

The diagram above is the one you will see most often at this level, but it is slightly misleading. Apart from the fact that some of the atoms don't seem to be forming enough bonds, it is all too neat and tidy!

There is a useful bit of YouTube video which shows a much bigger section of the structure. Nobody is suggesting that you should learn this or try to draw it, but it is interesting to see a more realistic part of the structure.

The physical properties of diamond

Diamond

• sublimes at ordinary pressures at almost 4000°C. Very strong carbon-carbon covalent bonds have to be broken throughout the structure before sublimation occurs.

• is very hard. This is again due to the need to break very strong covalent bonds operating in 3-dimensions.

• doesn't conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move.

• is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.

The giant covalent structure of graphite

The photo shows lumps of graphite and the marks they will make on paper.

Graphite is another allotrope of carbon. Allotropes are different forms of the same element in the same physical state. So carbon and diamond are both allotropes of carbon, because they are both chemically carbon, both solids, but obviously different.

Graphite has a layer structure which 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.

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 other part of the diagram being 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. And there are huge numbers of layers in each stack.

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 neighbouring 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 covalent bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalised 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 delocalised 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 delocalised electrons in one sheet and those in the neighbouring 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 delocalised electrons. The forces of attraction between the layers are much weaker.

The physical properties of graphite

Graphite

• has a high sublimation point, similar to that of diamond. In order for graphite to subllime, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure.

• 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.

• has a lower density than diamond. This is because of the relatively large amount of space that is "wasted" between the sheets.

• 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.

• conducts electricity. The delocalised electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.

The giant covalent structure of graphene

I am simply flagging up the existence of graphene, but without going into detail. Graphene is just a single sheet of the graphite structure.

The photo comes from this Wikipedia page. If you want to follow this up, you will probably find that page has far more than you want to know about graphene, and is far too complicated. A Google search should find something more accessible.

The giant covalent structure of silicon dioxide, SiO₂

This is probably less likely to come up in exams at this level then diamond and graphite, and I include it only to show that there are giant covalent substances which contain something other than carbon.

Silicon dioxide (also known as silica) is an important component of various rocks.

There are various different crystal forms of silicon dioxide. The easiest one to remember and draw is based on the diamond structure.

Crystalline silicon has the same structure as diamond. To turn it into silicon dioxide, all you need to do is to modify the silicon structure by including some oxygen atoms.

Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Don't forget that this is just a tiny part of a giant structure extending on all 3 dimensions.

The physical properties of silicon dioxide

Silicon dioxide

• has a high melting point - varying depending on what the particular structure is (remember that the structure given is only one of several possible structures), but around 1700°C. Very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs.

• is hard. This is due to the need to break the very strong covalent bonds.

• doesn't conduct electricity. There aren't any delocalised electrons. All the electrons are held tightly between the atoms, and aren't free to move.

• is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant structure.