Giant Structures: Metallic & Giant Covalent

Metallic Bonding

Metals are made of giant lattices of positively charged metal ions held together by a "sea" of delocalised electrons.

When metal atoms pack together, each one loses its outer-shell electrons. These electrons are no longer attached to any single atom — they are free to move throughout the whole structure. What is left behind is a regular, repeating arrangement of positive ions.

The bonding is the strong electrostatic attraction between the positive metal ions and the negative delocalised electrons that flow between them.

Key terms

Metallic bond — the electrostatic attraction between positive metal ions and a sea of delocalised electrons.

Delocalised electrons — outer-shell electrons that are free to move through the whole metal lattice and are not bound to one atom.

Explaining the Properties of Metals

Metallic bonding explains every typical metal property — and in the exam you must link the property back to the structure.

Good conductors of electricity. The delocalised electrons are free to move. When a voltage is applied, these electrons drift through the lattice and carry the charge. This works whether the metal is solid or molten.

Good conductors of heat. The mobile electrons (and closely packed ions) transfer kinetic energy quickly through the structure.

Malleable (can be hammered into shape) and ductile (can be drawn into wires). The layers of ions can slide over one another into new positions. The sea of electrons moves with them, so the strong attraction is maintained and the metal does not shatter.

High melting and boiling points. The electrostatic attraction between the positive ions and the delocalised electrons is strong, so a large amount of energy is needed to break the lattice apart.

Exam tip

Many marks are lost by saying "metals conduct because electrons move". Be precise: metals conduct because they contain delocalised electrons that are free to move and carry charge through the lattice.

Watch out

Generally the more outer electrons each atom donates, the stronger the metallic bond. This is why magnesium (donates 2 electrons, 2+ ions) has a higher melting point than sodium (donates 1 electron, 1+ ion).

Giant Covalent (Macromolecular) Structures

A giant covalent structure contains a huge number of atoms joined by a continuous network of strong covalent bonds. There are no separate small molecules — the whole structure is effectively one giant molecule.

Because every bond in the network is a strong covalent bond, these substances all have very high melting and boiling points — melting requires breaking many strong covalent bonds, which takes a great deal of energy.

The three you must know are diamond, graphite (both forms of carbon) and silicon dioxide (silica, $SiO_2$).

Diamond

In diamond, every carbon atom is bonded to four other carbon atoms by strong covalent bonds, arranged tetrahedrally. This builds a rigid three-dimensional network.

Properties of diamond, explained from structure:

Very hard — the rigid network of strong covalent bonds in all directions resists being deformed. This is why diamond is used in cutting tools and drill tips.
Very high melting point — a huge amount of energy is needed to break the many strong covalent bonds.
Does not conduct electricity — all four outer electrons of each carbon are used in bonding, so there are no free (delocalised) electrons to carry charge.

Graphite

Graphite is also made only of carbon, but each carbon atom is bonded to only three other carbons. This forms flat layers of carbon atoms arranged in hexagonal rings.

The fourth outer electron of each atom is delocalised and free to move between the layers. The layers themselves are held to each other only by weak forces (weak intermolecular forces), not by covalent bonds.

Properties of graphite, explained from structure:

Soft and slippery — the weak forces between layers let the layers slide over each other easily. This makes graphite useful as a lubricant and in pencil "lead".
Conducts electricity — each carbon has one delocalised electron, and these are free to move, carrying charge through the layers. (Diamond does not, because it has no spare electrons.)
High melting point — within each layer the carbon atoms are still joined by strong covalent bonds, so melting needs a lot of energy.

Exam tip

A classic compare question: Why does graphite conduct but diamond does not? Answer: in graphite each carbon forms only three bonds, leaving one delocalised electron free to move; in diamond each carbon forms four bonds, leaving no free electrons.

Silicon Dioxide (Silica)

Silicon dioxide, $SiO_2$, found in sand and quartz, is another giant covalent structure. Each silicon atom is covalently bonded to oxygen atoms, and each oxygen to silicon, in a continuous network similar to diamond.

As a result it shares diamond's general properties: it is hard, has a very high melting point, and does not conduct electricity (no free electrons).

Comparing the Four Structure Types

You must be able to compare the four main types of structure and explain their properties from their bonding.

Watch out

Simple molecular substances have low melting points not because covalent bonds are weak — the bonds inside each molecule are strong. It is the forces between the molecules that are weak and easily broken. Never write that you "break the covalent bonds" when melting a simple molecular substance like iodine.

Real world

The same element, carbon, forms diamond (hardest natural material, used in drills) and graphite (soft, used in pencils and as a lubricant). Identical atoms — completely different properties — purely because the arrangement of atoms and bonds is different. Structure determines properties: that idea runs through the whole of this topic.