giant covalent substances

Higher and foundation tiers

The allotropes of carbon

What do diamonds and a pencil lead have in common? At first it may appear that they have very little in common. Diamond is the hardest substance known, it is an electrical insulator, they are semi-transparent and they sparkle! Graphite on the other hand is often used in pencil leads; it is a dull coloured material which is so soft it can be used to write with on paper. It is also slippery to the touch and is used as a lubricant in industry. It is also a good electrical and thermal conductor. It is also worth mentioning that diamond is also an excellent thermal conductor.

These materials seem so different from each other that you may think there are no similarities or possible connections between them. Remarkably they are both made from the same element. Graphite and diamond are both forms of the element, namely carbon. Obviously the way the carbon atoms combine to form these two substances must be very different to result in two very different substances. We call these different forms of the same element allotropes.

and diamond are not the only allotropes of carbon, there are also the fullerenes. Carbon is probably the most remarkable element in the periodic table, it for example forms more compounds than all the other elements put together! It forms small molecules all to way up to giant macromolecular structure and it is one of the main elements in many polymers.

The reason that diamond and graphite are so different is the way in which the carbon atoms arrange themselves in forming these substances. Carbon can form substances made up of only a few atoms joined together - to form small molecular substances. However carbon is not only capable of forming substances consisting of small molecules but they can also link together to form giant covalent structures. The carbon atoms in these giant "molecules" can be arranged in different ways and so produce different substances. Carbon is one of only a few elements that are able to form these giant covalent structures.


The image below shows an illustration of the structure of a diamond. Each of the black spheres represents an atom of carbon. Each atom of carbon makes 4 strong covalent bonds. This giant structure is often called a macromolecular giant structure or giant covalent structure. All the carbon atoms are arranged in a tetrahedral structure as shown in the image. The structure is clearly very ordered (or crystalline) and very symmetrical. All the covalent bonds in the structure are strong and being highly symmetrical means there are no weak points in the structure. This makes diamond very hard. Lots of strong covalent bonds and a giant structure also mean the melting point will be high. Diamond is an electrical insulator since all the electrons are held tightly in covalent bonds and are not free to move.

The structure of diamond. Diamond has a giant macromolecular structure.

Silica- Silicon dioxide

Silicon is underneath carbon in group four of the periodic table. You may notice lots of similarities between the structure of silicon dioxide (sand) shown below and that of diamond shown above. In the image below the green spheres are silicon atoms and the red spheres represent oxygen atoms. The silicon atoms form a tetrahedral structure with four oxygen atoms around a central silicon atom, this small tetrahedral shaped molecule is shown sitting in front of the giant covalent structure in the image above. This tetrahedral shape repeats to build up the giant silicon dioxide macromolecular structure. Silicon forms strong covalent bonds with oxygen, so silicon dioxide having a giant structure with lots of strong bonds has a high melting and boiling point. There are no free delocalised electrons present so it is an electrical insulator. The image below shows the structure of silica or silicon dioxide.

The structure of silica, silica has a giant macromolecular structure.


Graphite like diamond has a giant macromolecular structure but the carbon atoms in graphite are arranged in a very different way to that found in diamond. The image below shows how the carbon atoms are arranged in graphite. The carbon atoms in the image are coloured to help you visualize the structure of graphite. The carbon atoms are arranged in flat layers of hexagons, shown in green, blue and red in the image below. By studying the picture you can see that each carbon atom makes only three covalent bonds, unlike the four covalent bonds each carbon atom makes in diamond. Since each carbon atom should make four bonds, this means that there are free electrons within the structure. These free or delocalised electrons are found between the layers of hexagons and being delocalised and free to move means that graphite is a good electrical and thermal conductor.

The structure of graphite.  The graphite structure consists of flat planes of hexagons.

The layered structure of graphite.The flat layers of carbon atoms in the hexagon structure are held together by strong covalent bonds between the carbon atoms; however there are only very weak intermolecular bonds (Van der Waals bonds) between the separate layers of hexagons. This means that if a small force is applied to one layer it can easily separate and slide free from the other layers. This is one of the main reasons why graphite is soft. When you write with a pencil the force that you push down on the tip with is big enough to break the weak intermolecular bonds between the hexagon layers and leave behind a layer of flat hexagons on the paper. If you scribble on paper over and over again to leave a thick pencil mark then rub your finger over the thick mark it will feel very slippery. This is because the flat layers of hexagons left on the paper are able to slide over each other and act as a lubricant.

Key Points

Practice questions

Check your understanding - Questions on giant structures

Check your understanding - Additional questions on giant structures

Check your understanding - Quick Quiz on covalent giant structures.