Bonding, structure and properties of matter: covalent bonding and simple molecular substances

What you'll learn

This guide covers covalent bonding and simple molecular substances as specified in the AQA GCSE Chemistry syllabus. You'll learn how non-metal atoms share electrons to form molecules, how to represent these bonds using diagrams, and why simple molecular substances have characteristic physical properties. Understanding these concepts is essential for explaining the behaviour of common substances like water, methane and carbon dioxide.

Key terms and definitions

Covalent bond — a strong electrostatic attraction between a shared pair of electrons and the nuclei of two atoms

Simple molecular substance — a substance made of molecules containing a small number of atoms joined by covalent bonds, with weak intermolecular forces between molecules

Dot-and-cross diagram — a drawing showing the arrangement of outer shell electrons in atoms or molecules, using dots and crosses to distinguish electrons from different atoms

Intermolecular forces — weak forces of attraction between molecules (much weaker than covalent bonds)

Shared pair — a pair of electrons in a covalent bond, with one electron contributed by each atom

Lone pair — a pair of electrons in the outer shell that is not involved in bonding

Diatomic molecule — a molecule containing exactly two atoms (e.g. H₂, O₂, Cl₂)

Relative molecular mass (Mr) — the sum of the relative atomic masses of all atoms in a molecule

Core concepts

Formation of covalent bonds

Covalent bonding occurs between non-metal atoms. When two non-metal atoms approach each other, they share one or more pairs of electrons to achieve full outer electron shells, similar to the electron configuration of noble gases.

Key principles:

• Each shared pair of electrons creates one covalent bond
• Atoms share electrons rather than transferring them completely
• The shared electrons are attracted to the nuclei of both atoms, creating a strong electrostatic attraction
• Covalent bonds are very strong and require significant energy to break

Common bonding patterns follow from electron shell structure:

• Hydrogen (Group 1) forms 1 covalent bond to achieve 2 outer electrons
• Carbon (Group 4) forms 4 covalent bonds to achieve 8 outer electrons
• Nitrogen (Group 5) forms 3 covalent bonds to achieve 8 outer electrons
• Oxygen (Group 6) forms 2 covalent bonds to achieve 8 outer electrons
• Chlorine (Group 7) forms 1 covalent bond to achieve 8 outer electrons

Drawing dot-and-cross diagrams

Dot-and-cross diagrams show only the outer shell (valence) electrons. Different symbols (dots vs crosses) distinguish electrons from different atoms, making it easier to see which electrons are shared.

Step-by-step method:

1. Determine how many outer electrons each atom has (from its group number)
2. Work out how many bonds each atom needs to form
3. Draw circles or brackets representing each atom
4. Add outer electrons using dots for one atom, crosses for another
5. Position shared electrons in the overlap region between atoms
6. Check each atom has a full outer shell (8 electrons, or 2 for hydrogen)

Examples of common molecules:

Hydrogen (H₂): Each hydrogen atom contributes 1 electron to form a single covalent bond. Both atoms achieve 2 outer electrons.

Water (H₂O): Oxygen shares electrons with two hydrogen atoms, forming two single covalent bonds. Oxygen has 6 outer electrons, so needs 2 more. Each hydrogen needs 1 electron. The oxygen atom also has two lone pairs.

Methane (CH₄): Carbon shares electrons with four hydrogen atoms, forming four single covalent bonds. The carbon achieves 8 outer electrons; each hydrogen achieves 2.

Ammonia (NH₃): Nitrogen shares electrons with three hydrogen atoms, forming three single covalent bonds. Nitrogen also has one lone pair.

Carbon dioxide (CO₂): Carbon shares two pairs of electrons with each oxygen atom, forming two double covalent bonds. Each atom achieves 8 outer electrons.

Oxygen (O₂): Each oxygen atom shares two pairs of electrons, forming a double covalent bond. Each oxygen also has two lone pairs.

Nitrogen (N₂): Each nitrogen atom shares three pairs of electrons, forming a triple covalent bond. Each nitrogen also has one lone pair.

Structure of simple molecular substances

Simple molecular substances consist of small, discrete molecules. The atoms within each molecule are held together by strong covalent bonds, but the molecules themselves are held together only by weak intermolecular forces.

This structure has crucial implications:

• Individual molecules have fixed compositions (e.g. each water molecule is always H₂O)
• Molecules do not join to form giant structures
• The weak forces between molecules are easily overcome
• Simple molecular substances exist as gases, liquids or low-melting-point solids at room temperature

Common examples include:

• Elements: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂, P₄, S₈
• Compounds: H₂O, CH₄, NH₃, CO₂, HCl

Physical properties of simple molecular substances

The physical properties of simple molecular substances arise directly from their bonding and structure.

Melting and boiling points:

• Simple molecular substances have low melting and boiling points
• When these substances melt or boil, only the weak intermolecular forces are overcome
• The strong covalent bonds within molecules remain intact
• Less energy is required to overcome weak intermolecular forces than to break strong covalent bonds
• Therefore, melting and boiling points are relatively low

Example: Water boils at 100°C, whereas ionic sodium chloride boils at 1413°C. The intermolecular forces in water are much weaker than the ionic bonds in sodium chloride.

State at room temperature:

Simple molecular substances may be:

• Gases (e.g. O₂, N₂, CO₂, CH₄, Cl₂) — very weak intermolecular forces
• Liquids (e.g. H₂O, Br₂) — weak to moderate intermolecular forces
• Solids (e.g. I₂, S₈) — moderate intermolecular forces (but still much weaker than covalent bonds)

Trend in boiling points:

For similar molecules, boiling point generally increases with relative molecular mass (Mr) because larger molecules have stronger intermolecular forces.

Example: The halogens show this trend clearly:

• Fluorine (F₂, Mr = 38) — gas at room temperature
• Chlorine (Cl₂, Mr = 71) — gas at room temperature
• Bromine (Br₂, Mr = 160) — liquid at room temperature
• Iodine (I₂, Mr = 254) — solid at room temperature

Electrical conductivity:

• Simple molecular substances do not conduct electricity in any state
• They contain no free-moving charged particles (no delocalised electrons or ions)
• The electrons are all held in covalent bonds or as lone pairs
• This applies to both elements (like O₂) and compounds (like CH₄)

Solubility:

• Many simple molecular substances are insoluble in water (e.g. methane, oxygen, nitrogen)
• Some simple molecular substances dissolve in water without ionising (e.g. sugar, ethanol)
• A few simple molecular substances react with water to form ions (e.g. HCl forms H⁺ and Cl⁻ ions)
• Simple molecular substances are often soluble in organic solvents like hexane or cyclohexane

Comparing covalent bonding with other bonding types

Understanding the differences between covalent, ionic and metallic bonding is essential for explaining properties.

Covalent vs ionic bonding:

Important distinction:

The term "covalent bonding" describes the type of bond (shared electrons). "Simple molecular" describes the structure (small molecules with weak forces between them). Not all covalent substances are simple molecular — giant covalent structures like diamond and silicon dioxide also contain covalent bonds but have very different properties.

Representing molecules

At GCSE level, you must be able to represent molecules in three ways:

Displayed formula: Shows all atoms and all covalent bonds as lines. Example for methane: each C-H bond shown as a line.

Molecular formula: Shows the number of each type of atom. Example: CH₄, H₂O, CO₂

Dot-and-cross diagram: Shows outer shell electrons, distinguishing those from different atoms.

You should be able to:

• Draw dot-and-cross diagrams for simple molecules
• Interpret displayed formulae to identify bonding patterns
• Work out molecular formulae from diagrams
• Explain bonding in terms of electron sharing

Worked examples

Example 1: Drawing a dot-and-cross diagram for chlorine (Cl₂)

Question: Draw a dot-and-cross diagram for a molecule of chlorine. (2 marks)

Solution:

1. Chlorine is in Group 7, so each atom has 7 outer electrons
2. Each atom needs 1 more electron to complete its outer shell
3. The two chlorine atoms share one pair of electrons
4. Each atom should be shown with 8 outer electrons (7 from itself, 1 shared)

Correct diagram shows:

• Two chlorine atoms (circles or Cl symbols)
• 6 unshared electrons on each atom (3 lone pairs) — use dots for one atom, crosses for the other
• 2 electrons in the overlap/shared region (1 dot, 1 cross)

Mark scheme:

• 1 mark: correct number of outer electrons shown for both atoms
• 1 mark: one shared pair correctly shown between the atoms

Example 2: Explaining physical properties

Question: Methane has a boiling point of -161°C. Explain why methane has such a low boiling point. (3 marks)

Model answer:

Methane is a simple molecular substance (1 mark). It has weak intermolecular forces between the CH₄ molecules (1 mark). Only a small amount of energy is needed to overcome these weak forces, so the boiling point is low (1 mark).

Common errors to avoid:

• Do NOT say "weak covalent bonds" — the covalent bonds within molecules are strong
• Do NOT say "bonds between molecules" — there are no bonds between molecules, only weak intermolecular forces
• Do focus on the forces BETWEEN molecules, not within them

Example 3: Predicting properties from structure

Question: The table shows information about three substances.

(a) Which substance is a simple molecular substance? (1 mark)

(b) Explain your answer to part (a). (2 marks)

Solution:

(a) Both B (H₂O) and C (CH₄) are simple molecular substances (1 mark for either)

(b) Water/methane has a low melting point (1 mark), which indicates weak intermolecular forces between molecules that are easily overcome (1 mark).

Alternative approach: MgO has a very high melting point indicating it is ionic, not simple molecular. Simple molecular substances have low melting points.

Common mistakes and how to avoid them

• Confusing bonds within molecules with forces between molecules. Remember: strong covalent bonds hold atoms together within each molecule; weak intermolecular forces exist between separate molecules. When simple molecular substances melt or boil, intermolecular forces are overcome, NOT covalent bonds.

• Saying simple molecular substances have "weak covalent bonds." This is incorrect. The covalent bonds are strong. The intermolecular forces are weak. Always specify what is weak (the forces between molecules).

• Drawing incorrect dot-and-cross diagrams by showing inner shell electrons. Only show outer (valence) shell electrons. For example, oxygen has 8 electrons total, but only show the 6 outer electrons in your diagram.

• Forgetting to check that atoms have full outer shells in dot-and-cross diagrams. After drawing, count the electrons around each atom. Most atoms need 8 (except hydrogen, which needs 2). If the numbers are wrong, revise your diagram.

• Confusing simple molecular substances with giant covalent structures. Diamond and graphite contain covalent bonds but are NOT simple molecular — they're giant covalent structures with very high melting points. Simple molecular substances have discrete, small molecules.

• Stating that simple molecular substances conduct electricity when dissolved. Most don't. Only substances that form ions when they dissolve (like HCl) conduct electricity. Sugar dissolves but doesn't conduct because it remains as molecules.

Exam technique for "Bonding, structure and properties of matter: covalent bonding and simple molecular substances"

• Command word "Explain" requires you to give reasons or mechanisms. For properties questions, link structure → bonding/forces → property. Example: "low melting point because weak intermolecular forces which require little energy to overcome" (3 marking points).

• Dot-and-cross diagram questions: Show only outer electrons, use two different symbols consistently, ensure shared electrons are clearly in the overlap region. Typically worth 2 marks: 1 for correct total electrons, 1 for correct sharing/arrangement.

• "State" or "Give" questions need short, factual answers without explanation. Example: "State why oxygen does not conduct electricity" → "no free-moving charged particles/electrons" (1 mark). Don't waste time explaining in detail.

• 6-mark questions on bonding often require comparisons between bonding types. Structure your answer with clear paragraphs for each type, using scientific vocabulary (covalent bond, intermolecular forces, ionic lattice, delocalised electrons). Link structure to properties explicitly.

Quick revision summary

Covalent bonds form when non-metal atoms share pairs of electrons, creating strong electrostatic attractions. Simple molecular substances contain small molecules with strong covalent bonds within molecules but weak intermolecular forces between molecules. This gives them characteristic low melting and boiling points. They don't conduct electricity because they lack free-moving charged particles. Dot-and-cross diagrams show outer shell electrons, with shared pairs in overlap regions. Physical properties depend on the strength of intermolecular forces, which increase with molecular size.