Types of substance and their properties

Substances that consist of giant covalent structures are solids with very high melting points. All of the atoms in these structures are linked to other atoms by strong covalent bonds. These bonds must be overcome to melt or boil these substances.

Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures.

That is the difference from a small molecule in one line: here there is nothing weak to overcome. Melting means breaking covalent bonds themselves, which is why the melting points are measured in thousands of degrees.

You should be able to recognise giant covalent structures from diagrams showing their bonding and structure — every atom is joined by lines to its neighbours, and the bonds run off the edge of the picture rather than stopping.

In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure, so diamond is very hard, has a very high melting point and does not conduct electricity.

You should be able to explain the properties of diamond in terms of its structure and bonding — which means giving the reason, not repeating the property.

In graphite, each carbon atom forms three covalent bonds with three other carbon atoms, forming layers of hexagonal rings which have no covalent bonds between the layers.

In graphite, one electron from each carbon atom is delocalised.

The whole of graphite's behaviour follows from those two sentences. Three bonds instead of four leaves one outer electron spare — and that spare electron is what makes graphite conduct.

You should be able to explain the properties of graphite in terms of its structure and bonding, and you should know that graphite is similar to metals in that it has delocalised electrons. That comparison is worth holding on to: graphite is a non-metal that conducts, and it conducts for exactly the reason a metal does.

Graphene is a single layer of graphite and has properties that make it useful in electronics and composites.

You should be able to explain the properties of graphene in terms of its structure and bonding. Because it is one layer of graphite, the same explanations apply: each carbon atom is covalently bonded to three others, which makes the sheet very strong for its mass, and one electron per atom is delocalised, which makes it an excellent conductor. Being one atom thick, it is also almost transparent — which is what makes it useful in electronics, and why adding it to another material gives a composite that is stronger without being much heavier.

Fullerenes are molecules of carbon atoms with hollow shapes. The structure of fullerenes is based on hexagonal rings of carbon atoms, but they may also contain rings with five or seven carbon atoms. The first fullerene to be discovered was Buckminsterfullerene (C60), which has a spherical shape.

Carbon nanotubes are cylindrical fullerenes with very high length to diameter ratios. Their properties make them useful for nanotechnology, electronics and materials.

You should be able to recognise graphene and fullerenes from diagrams and descriptions of their bonding and structure, and give examples of the uses of fullerenes, including carbon nanotubes.

Uses worth naming: fullerene cages can carry a drug molecule inside them for delivery in medicine; fullerenes have a very large surface area, which makes them useful as catalysts or catalyst supports; and carbon nanotubes are used to reinforce materials such as tennis racquets and aircraft parts, and in electronics, because they are strong along their length and conduct.

Ionic compounds have regular structures — giant ionic lattices — in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions.

These compounds have high melting points and high boiling points because of the large amounts of energy needed to break the many strong bonds.

When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and so charge can flow.

Knowledge of the structures of specific ionic compounds other than sodium chloride is not required.

Substances that consist of small molecules are usually gases or liquids that have relatively low melting points and boiling points.

These substances have only weak forces between the molecules — intermolecular forces. It is these intermolecular forces that are overcome, not the covalent bonds, when the substance melts or boils.

The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points.

These substances do not conduct electricity because the molecules do not have an overall electric charge.

You should be able to use the idea that intermolecular forces are weak compared with covalent bonds to explain the bulk properties of molecular substances. Say both halves: the covalent bonds inside are strong, and the forces between the molecules are weak. Only the second half is overcome on melting or boiling, and that is why the melting points are low.

Polymers have very large molecules. The atoms in the polymer molecules are linked to other atoms by strong covalent bonds. The intermolecular forces between polymer molecules are relatively strong, and so these substances are solids at room temperature.

A polymer sits between the small molecules and the giant covalent structures. It is still made of separate molecules, so the forces between them can still be overcome — but the molecules are so long that those forces add up along the whole chain, which is why a polymer is a solid where a small molecule of the same elements would be a gas.

You should be able to recognise polymers from diagrams showing their bonding and structure — look for a repeating unit, drawn in brackets, with the bonds continuing out of each side.

Metals consist of giant structures of atoms arranged in a regular pattern.

The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure. The sharing of delocalised electrons gives rise to strong metallic bonds.

You should be able to recognise substances as metallic giant structures from diagrams showing their bonding. The give-away is a regular grid of identical positive circles with loose electrons drawn between them — no negative ions, and no separate molecules.

Because the delocalised electrons can move, metallic bonding will explain, in the next sub-topic, why metals conduct electricity and heat, why they have high melting points, and why they can be bent and shaped.

Metals have giant structures of atoms with strong metallic bonding. This means that most metals have high melting and boiling points.

In pure metals, atoms are arranged in layers, which allows metals to be bent and shaped. Pure metals are too soft for many uses and so are mixed with other metals to make alloys, which are harder.

You should be able to explain why alloys are harder than pure metals in terms of distortion of the layers of atoms in the structure of a pure metal.

The explanation runs like this. In a pure metal the atoms are all the same size and sit in regular layers, so the layers can slide over one another and the metal bends easily. An alloy contains atoms of a different size, which disrupt the regular arrangement and distort the layers. The layers can no longer slide over each other, so the alloy is harder.

Metals are good conductors of electricity because the delocalised electrons in the metal carry electrical charge through the metal.

Metals are good conductors of thermal energy because energy is transferred by the delocalised electrons.

Both answers are the same sentence with one word changed, and both depend on the delocalised electrons being free to move through the whole structure — which is exactly what metallic bonding provides.

Every diagram of bonding you are asked to draw is a model, and every model leaves something out. Questions ask what — so learn each model with its limitation attached, and answer with the specific thing missing rather than "it is not accurate".

The same caution applies to a drawing of a giant structure. A diagram of an ionic lattice or of diamond can only show a few dozen ions or atoms, and the real structure continues in every direction for millions of them. That is why answers about melting points must talk about the many strong bonds throughout the whole structure, not about the handful drawn on the page.