OxfordAQA International GCSE Chemistry: How bonding and structure are related to the properties of substances (9202)

This study guide teaches section 3.2.2 of OxfordAQA’s International GCSE Chemistry (9202) specification (Version 6.2; it applies to exams May/June 2018 onwards): how bonding and structure are related to the properties of substances. It works through statements a to g in order. The qualification is untiered, so every statement here applies to every candidate, and section 3.2.2 has no required practical of its own.

The whole section rests on one idea. To melt or boil a substance, you must overcome the forces that hold its particles in place. To conduct electricity, a substance must contain charged particles that are free to move. If you can name the particles in a structure and the forces between them, you can explain every property in this section.

A free ten-minute diagnostic will show which parts of the course need attention.

Coverage of section 3.2.2

Statement	Skills you need
3.2.2 a	Describe a giant ionic lattice and explain the high melting and boiling points of ionic compounds
3.2.2 b	Explain why ionic compounds conduct when melted or dissolved in water
3.2.2 c	State that simple molecular substances have relatively low melting and boiling points
3.2.2 d	Explain that weak intermolecular forces, not covalent bonds, are overcome on melting or boiling
3.2.2 e	Explain why simple molecular substances do not conduct electricity
3.2.2 f	Describe giant covalent structures (diamond, graphite, silicon dioxide) and explain their very high melting points; recognise other giant structures from diagrams
3.2.2 g	Explain why metals conduct heat and electricity

How the bonds form belongs to section 3.2.1. Electron arrangements are covered in our guide to atoms and the periodic table for this course.

Four structure types at a glance

Structure	Particles	What holds them together	Melting point	Conducts?
Giant ionic lattice	Positive and negative ions	Strong electrostatic attraction in all directions	High	Only when melted or dissolved
Simple molecules	Small molecules	Strong covalent bonds inside each molecule; weak intermolecular forces between molecules	Relatively low	No
Giant covalent	Atoms	Strong covalent bonds throughout the whole structure	Very high	No (graphite is the exception)
Metallic	Positive ions and delocalised electrons	Strong electrostatic attraction between the ions and the electrons	Usually high	Yes, as solid and liquid

Ionic compounds (3.2.2 a and b)

The giant ionic lattice

An ionic compound such as sodium chloride is a giant ionic lattice: a regular, repeating three-dimensional arrangement of oppositely charged ions. In sodium chloride, each sodium ion (Na⁺) is surrounded by six chloride ions (Cl⁻), and each chloride ion is surrounded by six sodium ions. The pattern continues in every direction until the crystal ends. There are no separate “NaCl molecules”; the formula only gives the ratio of ions, 1 : 1.

The forces between oppositely charged ions are strong electrostatic forces of attraction acting in all directions through the lattice: every ion is attracted to all of its neighbours.

Sodium chloride is the only ionic structure the specification requires you to know.

Why the melting point is high

To melt an ionic solid, the ions must be pulled out of their fixed positions so that they can move past one another. That means overcoming a very large number of strong electrostatic attractions. A large amount of energy is needed, so the melting point and boiling point are high. Sodium chloride, for example, melts at about 801 °C.

A full answer has three linked parts:

1. Structure: giant ionic lattice of oppositely charged ions.
2. Force: strong electrostatic attraction acting in all directions; there are many of these.
3. Energy: overcoming them takes a lot of energy, so the melting point is high.

Why ionic compounds conduct only when melted or dissolved

An electric current is a flow of charge. In an ionic compound, the charge carriers are the ions.

• Solid: each ion is locked into its place in the lattice. It vibrates on the spot but never travels, so no current flows.
• Melted (molten): the lattice has broken down and the ions are free to move. They carry the current, so the liquid conducts.
• Dissolved in water: the lattice comes apart and the ions mix with the water molecules. Now mobile, they carry the current through the solution.

Worked example 1: explaining a conductivity result

A student tests solid potassium bromide with a simple circuit. The bulb does not light. She heats the solid until it melts, and the bulb lights. Explain both observations.

• Potassium bromide is an ionic compound made of K⁺ and Br⁻ ions in a giant lattice.
• In the solid, strong electrostatic forces lock each ion in place, so no charge can travel. The bulb stays off.
• When the solid melts, the ions are free to move. Moving ions carry the current, so the bulb lights.

Simple molecular substances (3.2.2 c, d and e)

Low melting and boiling points

Substances made of small molecules, such as H₂, O₂, N₂, Cl₂, HCl, H₂O, NH₃ and CH₄, have relatively low melting and boiling points. At room temperature they may be gases, liquids or solids that melt easily.

Two kinds of force

A simple molecular substance contains two very different kinds of force:

• Covalent bonds inside each molecule. These are strong. They hold the atoms of one molecule together.
• Intermolecular forces between neighbouring molecules. These are weak compared with covalent bonds.

When the substance melts or boils, the molecules move further apart or start to move freely. Only the intermolecular forces are overcome. The covalent bonds are not broken, so the molecules stay whole. When water boils, the steam is still made of H₂O molecules; it is not hydrogen and oxygen.

Weak forces need only a small amount of energy to overcome, so the melting and boiling points are relatively low.

Worked example 2: comparing the energies

Use these values for hydrogen chloride. Energy needed to boil 1 mol of liquid HCl: 16 kJ. Energy needed to break the H–Cl bonds in 1 mol of HCl molecules: 432 kJ. How many times larger is the bond-breaking energy? What does this tell you about boiling?

• Each HCl molecule has one H–Cl bond, so the bond-breaking energy for 1 mol is 432 kJ.
• Ratio: 432 ÷ 16 = 27, so breaking the covalent bonds would need 27 times as much energy as boiling.
• Boiling therefore cannot involve breaking the covalent bonds. It only overcomes the much weaker intermolecular forces, which is why the boiling point is low.

Why simple molecules do not conduct

Molecules do not have an overall electric charge. There are no ions and no free electrons, so there are no charged particles to carry a current. This is true whether the substance is solid, liquid or gas.

Worked example 3: deciding states from data

Use these values. At 25 °C, state whether each substance is a solid, liquid or gas.

Substance	Melting point / °C	Boiling point / °C
A	−114	−85
B	−23	77
C	5	80
D	842	1484

Method: compare 25 °C with each value. Below the melting point means solid. Between the melting point and the boiling point means liquid. Above the boiling point means gas.

• A: 25 °C is above −85 °C, so gas.
• B: −23 < 25 < 77, so liquid.
• C: 5 < 25 < 80, so liquid. It stays liquid over a 75 °C range, from 5 °C to 80 °C.
• D: 25 °C is below 842 °C, so solid.

A, B and C could be simple molecular. D could be ionic, metallic or giant covalent; conductivity data would decide.

Giant covalent structures (3.2.2 f)

Atoms that share electrons do not always form small molecules. They can also form giant covalent structures, also called macromolecules or giant covalent lattices. In these, every atom is linked to other atoms by strong covalent bonds, and the network continues through the whole structure. There are no separate molecules and so no weak intermolecular forces to overcome.

The examples you must know are:

• Diamond: a form of carbon in which each carbon atom is bonded to four others.
• Graphite: a form of carbon in which each carbon atom is bonded to three others in layers.
• Silicon dioxide (silica): a three-dimensional network in which every Si atom has four O neighbours and every O atom bridges two Si atoms.

To melt any of these, you must break a huge number of strong covalent bonds. That needs a very large amount of energy, so they have very high melting points. Silicon dioxide, for example, melts at roughly 1700 °C. The detailed structures and other properties of diamond and graphite, such as hardness and graphite’s conductivity, belong to section 3.2.3.

Recognising a giant structure from a diagram

For an unfamiliar structure, ask:

1. Are the particles charged? Labels such as + and − on alternating particles suggest an ionic lattice.
2. Is there a sea of electrons? Positive ions surrounded by “e⁻” or dots labelled delocalised electrons show a metal.
3. Do the bonds stop? If the diagram shows small, separate groups of atoms with gaps between them, it is simple molecular.
4. Do the bonds continue? If lines (covalent bonds) join atoms to atoms with no end, often drawn with broken bonds at the edges to show the pattern carries on, it is giant covalent.

Worked example 4: classifying from a description

Diagram 1 shows boron and nitrogen atoms. Each boron atom is joined by lines to four nitrogen atoms, and each nitrogen atom to four boron atoms. The pattern runs off every edge of the diagram. Diagram 2 shows rings of eight sulfur atoms; each ring is separate from the next. Classify each structure and predict which has the higher melting point.

• Diagram 1: atoms only, joined by covalent bonds that continue in all directions. Giant covalent.
• Diagram 2: separate rings, each a molecule. Simple molecular. Melting overcomes only the weak forces between rings.
• Diagram 1 has the much higher melting point, because melting it requires breaking many strong covalent bonds.

Metals (3.2.2 g)

Section 3.2.1 describes a metal as a giant structure of positive ions with delocalised electrons moving between them. Those delocalised electrons explain conduction.

• Electrical conduction: when a voltage is applied, the delocalised electrons move through the whole structure, carrying charge. Metals conduct as solids and as liquids.
• Thermal conduction: the delocalised electrons move freely and carry energy quickly through the metal, so heat passes along it.

The key phrase is that conduction depends on the ability of electrons to move through the metal. In a metal, the charge carriers are electrons. In a molten ionic compound, they are ions. Keep the two apart.

Worked example 5: metal versus ionic solid

Solid magnesium conducts electricity but solid magnesium chloride does not. Explain.

• Magnesium contains delocalised electrons, which drift through the metal and carry charge.
• Magnesium chloride is ionic. In the solid, its ions are fixed in the lattice and cannot move, and it has no delocalised electrons. So there are no mobile charged particles and it does not conduct.

Common errors

• “Ionic bonds are broken when salt dissolves, so electrons flow.” In solution the charge carriers are ions, not electrons.
• Blaming weak covalent bonds for water’s low boiling point. The O–H bonds are strong; it is the forces between H₂O molecules that are weak.
• Calling silicon dioxide molecular. It is a giant covalent structure with no separate molecules.
• Saying giant covalent substances have “strong intermolecular forces”. There are no intermolecular forces to talk about; there are strong covalent bonds throughout.
• Leaving out the energy step. State that overcoming many strong forces takes a lot of energy.

Where to go next

• Practice questions with worked answers
• Revision notes for the final weeks
• Exam preparation for this course
• Every unit of this course, on the hub and the tick-off checklist

Official syllabus

Published by OxfordAQA: the International GCSE Chemistry (9202) specification, Version 6.2, valid for exams May/June 2018 onwards. This page follows its section 3.2.2, “How bonding and structure are related to the properties of substances”.