Giant Covalent Structures, Metals and Alloys

A giant covalent structure is one huge network in which a very large number of atoms are all joined to their neighbours by strong covalent bonds (shared pairs of electrons). There are no small separate molecules and no weak forces between molecules to overcome, because the whole solid is effectively a single connected lattice.

This is why giant covalent structures behave so differently from small molecules such as water or carbon dioxide. In a small molecule the covalent bonds inside the molecule are strong, but the forces between molecules are weak, so the substance melts easily. In a giant covalent structure there are no separate molecules at all, so melting means breaking the strong covalent bonds themselves.

The three examples you must be able to recognise from a diagram are:

Giant covalent structure	Made from	Everyday context
Diamond	Carbon atoms	Cutting tools, jewellery
Silicon dioxide (silica)	Silicon and oxygen atoms	Sand, quartz, glass
Graphite	Carbon atoms (layered)	Pencils, electrodes, lubricant

A giant covalent structure has no separate molecules. Every atom is locked into the network by covalent bonds, so the whole solid melts as one.

Giant covalent structures are solids with very high melting and boiling points. The reason is the same for all of them, and it earns the marks when you state it precisely.

To melt or boil a giant covalent structure you must break the strong covalent bonds that hold the atoms together. There are an enormous number of these bonds throughout the network, and each one is strong, so breaking enough of them to let the structure flow needs a very large amount of energy. That energy is only supplied at very high temperatures.

Compare the melting points to see the effect of bonding:

Substance	Structure type	Melting point (°C)	What must be broken to melt
Diamond	Giant covalent	About 3550	Strong covalent bonds
Silicon dioxide	Giant covalent	About 1710	Strong covalent bonds
Iodine	Small molecule	114	Weak intermolecular forces

Iodine is made of small $I_2$ molecules held together by weak intermolecular forces, so it melts far below diamond even though the bond inside each iodine molecule is covalent. The lesson: what you break when a substance melts is set by its structure, not just by whether covalent bonds are present.

A metal is also a giant structure, but it is held together by metallic bonding rather than covalent bonds. The outer-shell electrons of the metal atoms become delocalised, meaning they are no longer attached to any one atom and are free to move throughout the whole structure. This leaves the atoms as positive ions arranged in a regular pattern, surrounded by a "sea" of delocalised electrons.

The metallic bond is the strong electrostatic attraction between the positive ions and the sea of delocalised electrons. Because this attraction acts throughout the giant structure, metals have high melting and boiling points: a large amount of energy is needed to overcome the metallic bonding and separate the ions.

The positive ions and the shared electrons stay locked in this arrangement, which is why a metal is a strong, dense solid at room temperature.

The atoms in a pure metal are all the same size and are arranged in layers. Because the layers are made of identical atoms, they can slide over one another when a force is applied, without the metallic bonding breaking apart. This is why pure metals are relatively soft and can be bent, hammered and drawn into shapes.

This property is called being malleable (can be hammered into shape) and ductile (can be pulled into wires). Both come from the same cause: layers of atoms sliding past each other.

Pure metals are often too soft for a job that needs strength or hardness, such as a bridge girder or a knife blade. To make them harder, they are turned into alloys.

The regular layers are what make a pure metal easy to shape, and, as the next slide shows, disrupting those layers is exactly how an alloy is made harder.

An alloy is a mixture of a metal with one or more other elements. Alloys are harder than the pure metal they are based on, and the reason is exactly what examiners want stated in full.

The added atoms are usually a different size from the original metal atoms. When they sit in the structure they distort the layers, so the layers can no longer slide over each other easily. Because sliding is what let the pure metal deform, disrupting it makes the alloy harder and more rigid.

A full mark-scheme answer to "explain why alloys are harder than pure metals" contains two linked ideas:

1. The alloy contains atoms of different sizes.
2. These distort the layers so they cannot slide over each other as easily.

State both and the marks follow.

Metals are good conductors of electricity and good conductors of thermal energy (heat). Both properties come from the same feature: the delocalised electrons that are free to move through the whole structure.

Electrical conduction. When a metal is connected in a circuit, the delocalised electrons move through the structure and carry electrical charge from one end to the other. A material can only conduct electricity if it contains charged particles that are free to move; in a metal those particles are the delocalised electrons.

Thermal conduction. When one part of a metal is heated, the delocalised electrons there gain kinetic energy. Because they move freely, they transfer that energy quickly through the structure to cooler parts, so the whole piece of metal heats up. This is why a metal spoon in a hot drink soon feels hot all over.

Property	Cause in a metal
Conducts electricity	Delocalised electrons move and carry charge
Conducts thermal energy	Delocalised electrons transfer energy through structure
High melting/boiling point	Strong metallic bonding throughout the giant structure
Malleable and ductile	Layers of atoms slide over each other

The same delocalised electrons explain three key metal properties at once: electrical conduction, thermal conduction and (with the layers) shaping. Link the property back to the electrons and you have the explanation.

1. Saying a giant covalent structure melts by breaking "intermolecular forces"

Giant covalent structures have no separate molecules, so there are no intermolecular forces to break. Melting them means breaking the strong covalent bonds in the network. Only small molecular substances melt by overcoming weak intermolecular forces.

2. Giving a half-answer for why alloys are harder

"Alloys have different atoms" alone is not enough. You must say the different-sized atoms distort the layers so they cannot slide over each other. Both parts are needed for full marks.

3. Saying metals conduct because "electrons carry the current" without naming delocalised electrons

The mark scheme wants delocalised (or free) electrons that are free to move and carry charge. Just writing "electrons" without saying they are delocalised and mobile often does not gain the mark.

4. Confusing the causes of the two conduction properties

Both electrical and thermal conduction in a metal are due to delocalised electrons. Do not attribute thermal conduction to "vibrating ions" as the main GCSE reason; the required answer is that delocalised electrons transfer the energy.

5. Assuming an alloy is a compound

An alloy is a mixture, not a compound. The metal and the added element are mixed together, not chemically bonded into a new fixed-ratio substance.