GCSE Chemistry Bonding, structure and properties - Revision Guide, Questions and Exam Prep

GCSE Chemistry Bonding, Structure and Properties: Complete Paper 1 Revision Guide

GCSE Chemistry Bonding, Structure and Properties is one of the highest-frequency Paper 1 topics because it explains why substances behave so differently from one another. Examiners use this topic to test whether students can connect atomic structure to observable physical properties such as melting point, electrical conductivity, hardness and solubility. A student who can explain ionic, covalent and metallic bonding clearly — and link each type of bonding to its resulting structure and properties — will find a large number of Chemistry questions across both papers become significantly more manageable.

This topic builds directly on atomic structure and the periodic table, where electronic structure and the arrangement of outer-shell electrons are introduced. The type of bonding an element or compound forms is determined almost entirely by whether electrons are transferred, shared or delocalised — and that decision depends on the electronic structures of the atoms involved. Securing this link makes both topics stronger.

Ionic Bonding: Transfer of Electrons

Ionic bonding occurs when electrons are transferred from a metal atom to a non-metal atom. The metal loses one or more electrons to form a positively charged ion (cation). The non-metal gains those electrons to form a negatively charged ion (anion). Both ions then have full outer electron shells, which is the stable arrangement. The oppositely charged ions attract each other through strong electrostatic forces and arrange themselves into a regular repeating pattern called a giant ionic lattice.

The properties of ionic compounds follow directly from this structure:

• High melting and boiling points — a large amount of energy is needed to overcome the many strong electrostatic attractions throughout the entire lattice.


• Conduct electricity when dissolved in water or when molten — in these states, the ions are free to move and carry charge. In the solid state, ions are held in fixed positions and cannot move, so solid ionic compounds do not conduct.


• Brittle — when a force is applied, layers of ions shift so that ions of the same charge come alongside each other, causing repulsion and the lattice to shatter.


• Often soluble in water — water molecules can attract and surround the ions, pulling them away from the lattice.

In exam questions, the most common error is stating that ionic compounds "have ions that can move" without specifying that this only applies when molten or dissolved. Always include the state condition when discussing conductivity in ionic compounds.

Covalent Bonding: Sharing of Electrons

Covalent bonding occurs between non-metal atoms. Rather than transferring electrons, atoms share pairs of electrons. Each shared pair forms one covalent bond. Sharing electrons allows both atoms to achieve full outer shells without either atom losing electrons entirely. Covalent bonds are strong, but the properties of a covalently bonded substance depend critically on whether it has a simple molecular structure or a giant covalent structure.

Simple molecular substances such as water, carbon dioxide, chlorine and ammonia consist of small molecules held together by strong covalent bonds within each molecule. However, between molecules, the forces of attraction are weak intermolecular forces. These weak forces require very little energy to overcome, which is why simple molecular substances typically have low melting and boiling points and are often gases or liquids at room temperature. They do not conduct electricity because there are no free electrons or ions.

A very common mistake is confusing the strong covalent bonds within a molecule with the weak intermolecular forces between molecules. When a simple molecular substance melts or boils, it is the weak intermolecular forces that are broken — not the covalent bonds. Covalent bonds within the molecule remain intact. Writing "bonds break when it melts" without specifying which type of bond or force is one of the most reliable ways to lose marks in this topic.

Giant Covalent Structures

Some covalently bonded substances do not form small molecules but instead form giant structures in which every atom is bonded to multiple other atoms by strong covalent bonds throughout the entire structure. These are called giant covalent structures or macromolecular structures. Because breaking down these substances requires breaking vast numbers of strong covalent bonds, they have very high melting points. The most important examples are diamond, graphite, graphene and silicon dioxide.

Diamond: Each carbon atom forms four covalent bonds arranged tetrahedrally to four other carbon atoms. This produces an extremely rigid, three-dimensional network with no free electrons and no weak points. Diamond is the hardest natural substance, has a very high melting point, and does not conduct electricity because all electrons are used in bonding and none are free to move.

Graphite: Each carbon atom forms three covalent bonds to three other carbon atoms in flat hexagonal layers. The fourth electron from each carbon atom is delocalised and free to move between the layers. This gives graphite two unusual properties for a non-metal: it conducts electricity (because of the delocalised electrons) and it is soft and slippery (because the layers are held together only by weak intermolecular forces and can slide over each other). Graphite is used as an electrode and as a lubricant — both uses depend on these specific structural features.

Graphene: A single layer of graphite — one atom thick, arranged in a hexagonal lattice. Graphene is extremely strong, very light and conducts electricity. It is one of the thinnest and strongest materials known and is of significant interest in materials science and electronics.

Fullerenes: Molecules of carbon atoms arranged in hollow spheres or tubes. The most well-known is buckminsterfullerene (C₆₀), which consists of 60 carbon atoms arranged in pentagons and hexagons. Fullerenes have potential applications in drug delivery, nanotechnology and lubrication.

Silicon dioxide (SiO₂): Each silicon atom is covalently bonded to four oxygen atoms and each oxygen atom bridges two silicon atoms, forming a giant three-dimensional network similar to diamond. Silicon dioxide has a very high melting point and does not conduct electricity.

Metallic Bonding: Delocalised Electrons

In metals, the outer-shell electrons leave their parent atoms and become delocalised — free to move throughout the entire metallic structure. This leaves behind a regular lattice of positive metal ions. The strong electrostatic attraction between the positive ion lattice and the sea of delocalised negative electrons holds the metal together. This is metallic bonding.

The properties of metals follow directly from this structure:

• Good conductors of electricity — delocalised electrons can move freely through the structure and carry charge.


• Good conductors of heat — delocalised electrons transfer energy quickly through the metal.


• High melting points — many strong attractions between the positive ions and the delocalised electron sea must be overcome. (Though melting points vary considerably across metals.)


• Malleable and ductile — layers of positive ions can slide over each other without breaking the metallic bonds, because the delocalised electrons can rearrange around them. This allows metals to be hammered into shapes or drawn into wires.

Alloys — mixtures of metals, or metals with small amounts of non-metals — are often harder than pure metals. This is because differently sized atoms disrupt the regular lattice arrangement, making it harder for layers to slide over each other.

Comparing Structures: The Exam Strategy That Works

When a question asks you to compare the properties of different substances, use a consistent three-step chain for each substance: identify the bonding type and structure → explain the forces that must be overcome → state the resulting property. This structure prevents answers from becoming lists of isolated facts and ensures that every property is explained rather than just named.

6-mark model answer — comparing ionic and simple covalent substances: Sodium chloride is an ionic compound with a giant ionic lattice. Strong electrostatic forces act throughout the lattice in all directions, so a large amount of energy is needed to separate the ions. This gives sodium chloride a high melting point. When molten or dissolved, the ions are free to move, so it conducts electricity. By contrast, water is a simple covalent molecule. The covalent bonds within each molecule are strong, but the forces between molecules are weak. Very little energy is needed to overcome these intermolecular forces, so water has a low boiling point relative to ionic compounds. Water does not conduct electricity because it has no free ions or delocalised electrons.

Worked Examples

Question 1: Why does graphite conduct electricity but diamond does not?

Model answer: In graphite, each carbon atom forms three covalent bonds, leaving one electron per carbon atom delocalised. These delocalised electrons can move through the structure and carry charge. In diamond, each carbon atom forms four covalent bonds, so all electrons are used in bonding. There are no free electrons to carry charge, so diamond does not conduct electricity.

Question 2: Why do metals have high melting points but are also malleable?

Model answer: Metals have a giant structure of positive ions surrounded by delocalised electrons. Strong electrostatic attractions between the ions and the electron sea require a large amount of energy to overcome, giving metals high melting points. Metals are malleable because the layers of positive ions can slide over each other. The delocalised electrons rearrange around them as they move, so the metallic bonding is maintained and the metal does not shatter.

Question 3: Why does sodium chloride not conduct electricity as a solid, but does conduct when dissolved in water?

Model answer: In solid sodium chloride, the ions are held in fixed positions in the giant ionic lattice and cannot move, so charge cannot be carried and it does not conduct. When dissolved in water, the ions separate and are free to move through the solution. These mobile ions carry charge and allow the solution to conduct electricity.

Common Mistakes to Avoid

• Confusing intermolecular forces with covalent bonds. When simple molecular substances melt, weak intermolecular forces break — not covalent bonds.


• Saying ionic compounds conduct in the solid state. They only conduct when molten or dissolved, when ions are free to move.


• Describing graphite as simply "having layers" without explaining the delocalised electrons. The conductivity and softness of graphite both require separate explanations.


• Naming the structure without linking it to the property. Every answer about properties must end with the specific property the structure explains.


• Treating all covalent substances the same. Simple molecular and giant covalent substances have completely different properties — do not apply the same explanation to both.

Use this topic alongside atomic structure and the periodic table to reinforce how outer-shell electrons determine bonding type, and alongside quantitative chemistry where formula mass calculations require understanding of ionic and molecular formulas. Bonding and structure also underpin questions in organic chemistry, where all carbon compounds are covalently bonded, and in chemical analysis and atmosphere, where properties of different substances are used to identify and separate them.