Covalent Network

These are giant molecular lattice structures. This implies that strong covalent bonding holds their atoms together in a highly regular extended network. The bonding between the atoms goes on and on in three dimensions. Melting requires the separation of the species comprising the soild state, and boiling the separation of the species comprising the liquid state. Because of the large amount of energy needed to break huge numbers of covalent bonds, all giant covalent network structures have high melting points and boiling points and are insoluble in water. Diamond, graphite (allotropes of carbon) and quartz (silicon(IV) oxide, SiO2) are examples.

Diamond and Graphite

Diamond and graphite are crystalline forms of the element carbon. The model below illustrates that diamond has an immensley strong, rigid structure with a very high melting point (3800 °C). Indeed, diamond is the hardest known natural substance. As well as in jewellery, it is used to tip drill bits for drilling through rock and in cutting glass.

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A carbon atom with covalent bonds towards the four corners of a tetrahedron...

In graphite each carbon atom is covalently bonded to three others in the same plane. The bond angle is 120°, and so the carbon atoms can form six-membered rings that link up to form planes or flat sheets of carbon atoms. A section of a single plane of graphite is shown below.

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Like the covalent bonding in diamond, the bonding in the planes is very strong. The melting point of graphite is very high. It is used in containers for molten metals and in rocket exhausts.

In diamond all four outer electrons of each carbon atom are 'localised' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current. In graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane. Each carbon atom contributes one electron to a delocalised system of electrons that is also a part of the chemical bonding. The decolcalised electrons are free to move throughout the plane. For this reason, graphite conducts electricity along the planes of carbon atoms, but does not conduct in a direction at right angles to the plane.

Other forms of carbon, such as charcoal, are well known. All of these consist of microcrystals of graphite but in which the planes of carbon atoms are randomly orientated. They can conduct electricity in any direction. Use is made of carbon in industrial electrodes and electric furnaces. You may also have heard of carbon fibre. Long fibres of graphite can be made and used with other substances to make a variety of objects, such as turbine blades for jet engines and tennis racquets.

Diamond, graphite and quartz can be examined in 3D below:

Chemis3D is made available by Didier Collomb.
Thanks also to those who make their 'molecules' available on the Internet for download.

Silicon(IV) oxide (silicon dioxide)

You are likely to familiar with both carbon dioxide (CO2) and silicon dioxide (SiO2). Under normal conditions, carbon dioxide is a gas but silicon dioxide is a high melting point (1710 °C) solid. Pure sand is pulverized quartz. In its soild state, carbon dioxide (known as 'dry ice') changes directly from a solid to a gas (sublimes) at -78 °C.

Although C and Si are both in Group 4 of the Periodic Table, the small size of the carbon atom makes it possible for carbon to form double bonds with oxygen atoms. This results in a discrete group of three covalently bonded atoms, that is, a small molecule. The formula, CO2, is a molecular formula.

Silicon being a larger atom than carbon can form only a single covalent bond with an oxygen atom. Each silicon atom can bond to four oxygen atoms, and this gives rise to a giant covalent network structure in which each Si is bonded to four oxygens and each O to two silicon atoms. Again, the bonding between the atoms goes on and on in three dimensions. This results in a 1:2 ratio between Si and O atoms, and so SiO2 is really an empirical formula.

There are three crystalline forms of SiO2. For one of these, the basic arrangement of Si atoms is rather like carbon in diamond, but with an oxygen atom between the silicon atoms. In quartz spiral structures are present. The 3D model above displays a fragment of quartz large enough to show all of the oxygen atoms (red) around 4 of the silicon atoms.

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