Ionic, covalent and metallic bonding explained simply

Chemical bonding is one of the most important topics in GCSE Chemistry. It explains why atoms join together and why different substances behave the way they do. There are three types of bonding you need to know: ionic, covalent and metallic.

The good news is that all three follow a simple idea. Atoms bond because they want a full outer shell of electrons. The difference between ionic, covalent and metallic bonding is just how the atoms achieve that goal.

Ionic bonding happens between metals and non-metals. The metal atom transfers one or more electrons to the non-metal atom. This means the metal loses electrons and the non-metal gains them.

When an atom loses electrons it becomes a positively charged ion (a cation). When an atom gains electrons it becomes a negatively charged ion (an anion). The opposite charges attract each other strongly, and that electrostatic attraction is the ionic bond.

Take sodium chloride as an example. Sodium (Na) has one electron in its outer shell. Chlorine (Cl) has seven electrons in its outer shell and needs one more to fill it. Sodium transfers its outer electron to chlorine. Sodium becomes Na⁺ and chlorine becomes Cl⁻. The two ions are held together by the attraction between their opposite charges.

Ionic compounds do not exist as individual pairs of ions. Instead, billions of positive and negative ions arrange themselves into a giant ionic lattice. Each ion is surrounded by ions of the opposite charge in a regular, repeating 3D pattern.

This lattice structure is the reason ionic compounds have the properties they do.

Ionic compounds have high melting and boiling points. This is because the electrostatic forces between oppositely charged ions are very strong, and a lot of energy is needed to overcome them.

Ionic compounds do not conduct electricity when solid. The ions are locked in fixed positions in the lattice and cannot move. However, when melted or dissolved in water, the ions are free to move and can carry an electric charge, so the substance does conduct.

Ionic compounds tend to be soluble in water. Water molecules are polar and can separate the ions from the lattice.

Covalent bonding happens between non-metal atoms. Instead of transferring electrons, the atoms share one or more pairs of electrons. Each shared pair of electrons is one covalent bond.

By sharing electrons, both atoms can achieve a full outer shell. A molecule of water (H₂O), for example, has two covalent bonds. The oxygen atom shares one pair of electrons with each of the two hydrogen atoms. A molecule of oxygen gas (O₂) has a double covalent bond – two shared pairs of electrons between the two oxygen atoms.

Covalent bonds are strong. The electrostatic attraction between the shared pair of electrons and the nuclei of both atoms holds the atoms together tightly.

Most covalent substances are simple molecules. A simple molecule is a small group of atoms held together by strong covalent bonds within the molecule. However, the forces between molecules (intermolecular forces) are weak.

This is why simple molecular substances like water, oxygen and carbon dioxide have low melting and boiling points. It does not take much energy to overcome the weak intermolecular forces. The covalent bonds within the molecules do not break when the substance melts or boils.

Simple molecular substances do not conduct electricity. They have no free electrons or ions to carry a charge.

Some covalent substances form giant covalent structures (also called macromolecules). In these, every atom is bonded to its neighbours by strong covalent bonds in a continuous network. There are no separate molecules.

Diamond and silicon dioxide are examples. Diamond has each carbon atom bonded to four other carbon atoms in a rigid tetrahedral structure. This makes it extremely hard and gives it a very high melting point.

Graphite is another giant covalent structure, but with a twist. Each carbon atom bonds to three others, forming layers of hexagonal rings. The fourth outer electron from each carbon is delocalised, which means it can move freely between the layers. This is why graphite conducts electricity – those delocalised electrons carry the charge. The layers are held together by weak intermolecular forces, so they can slide over each other, making graphite soft and slippery.

Metallic bonding occurs in metals and alloys. In a metal, the atoms are packed closely together in a regular arrangement. The outer electrons from each metal atom become delocalised – they are no longer attached to any one atom and can move freely throughout the whole structure.

This creates a lattice of positive metal ions surrounded by a sea of delocalised electrons. The strong electrostatic attraction between the positive ions and the negative delocalised electrons is the metallic bond.

Metals have high melting and boiling points because the metallic bonds (the attraction between positive ions and delocalised electrons) are strong and require a lot of energy to break.

Metals are good conductors of electricity. The delocalised electrons are free to move through the structure and carry an electrical charge. This also makes metals good conductors of heat, as the free electrons can transfer kinetic energy quickly.

Metals are malleable (can be bent or hammered into shape) and ductile (can be drawn into wires). When a force is applied, the layers of ions can slide over each other without the bonding breaking, because the delocalised electrons can shift to maintain the attraction.

Dot-and-cross diagrams are a way of showing how electrons are arranged in a bond. You draw one atom's electrons as dots and the other atom's electrons as crosses. This makes it easy to see which electrons have been transferred (in ionic bonding) or shared (in covalent bonding).

For ionic bonding, draw the electron transfer and then show the resulting ions in square brackets with their charges. For example, in NaCl, the dot-and-cross diagram shows sodium's single outer electron being transferred to chlorine, giving Na⁺ with an empty outer shell (showing the full shell beneath) and Cl⁻ with a full outer shell of eight electrons.

For covalent bonding, draw the overlapping outer shells with the shared pair of electrons in the overlap. For example, in H₂, each hydrogen contributes one electron to the shared pair, giving both atoms a full outer shell of two electrons.

You do not normally draw dot-and-cross diagrams for metallic bonding at GCSE. Instead, you describe or sketch the model of positive ions in a sea of delocalised electrons.