AQA GCSE Chemistry 8462: Key ideas – Study Guide

This study guide teaches section 4.11, Key ideas, of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 with GCSE exams from June 2018 (version 1.1). The specification says these ideas are embedded throughout the subject content and will be assessed across all papers, so they can turn up on Paper 1 or Paper 2. Section 4.11 has no tier split of its own. Some of the content you use to explain the ideas is Higher tier only, and it is labelled here. Exams run every May/June for the life of the specification.

Use it with the Key ideas revision notes and the Key ideas practice questions. The course hub is AQA GCSE Chemistry and the printable checklist lists every topic. For more detail on the first two ideas, see the atomic structure revision notes and the chemical bonds and ionic bonding revision notes.

What this section covers

Section 4.11 lists eight key ideas. They add no new facts: each runs through topics 4.1 to 4.10. You must recognise the idea inside a question and use it to build a linked explanation.

#	Key idea (spec wording, shortened)	Where it is taught in 8462
1	Matter is made of atoms; about 100 naturally occurring elements	4.1.1.1 atoms, elements and compounds; 4.1.1.4–4.1.1.6
2	Elements show periodic relationships in their properties	4.1.2.1 the periodic table; 4.1.2.4–4.1.2.6 Groups 0, 1, 7
3	Periodic properties are explained by atomic structure	4.1.1.7 electronic structure; 4.1.2.3 metals and non-metals
4	Atoms bond by transferring or sharing electrons	4.2.1.1–4.2.1.5 ionic, covalent and metallic bonding
5	Shapes of molecules and arrangement of giant structures decide behaviour	4.2.2 properties of substances; 4.2.3 carbon structures
6	There are barriers to reaction, so reactions occur at different rates	4.5.1.2 reaction profiles; 4.6.1.3 collision theory; 4.6.1.4 catalysts
7	Reactions happen in only three ways: proton transfer, electron transfer, electron sharing	4.4.2 acids; 4.4.1.4 and 4.4.3.5 redox (Higher tier only); 4.2.1.4 covalent bonds
8	Energy is conserved: it is neither created nor destroyed	4.5.1.1 exothermic and endothermic reactions

How the key ideas are assessed

Paper 1 covers topics 1–5 and Paper 2 covers topics 6–10. Each paper is a written exam of 1 hour 45 minutes, worth 100 marks and 50% of the GCSE, set at Foundation and Higher tier. Paper 2 questions may draw on fundamental concepts from sections 4.1 to 4.3. You are given a periodic table in the exam.

1. Atoms and elements

All substances are made of atoms. An atom is the smallest part of an element that can exist. There are about 100 different elements, each with its own chemical symbol (O for oxygen, Na for sodium). Compounds contain two or more elements chemically combined in fixed proportions, and can only be separated into elements by chemical reactions.

Atoms are tiny. The radius of an atom is about 0.1 nm (1 × 10⁻¹⁰ m). The radius of the nucleus is less than 1/10 000 of that of the atom (about 1 × 10⁻¹⁴ m). If a nucleus had a radius of 1 cm, the atom would have a radius of about 100 m. Almost all the mass is in the nucleus.

Particle	Relative charge	Relative mass
Proton	+1	1
Neutron	0	1
Electron	–1	Very small

The atomic number is the number of protons. The mass number is protons + neutrons. Isotopes are atoms of the same element with different numbers of neutrons.

Worked example 1. Find the numbers of protons, neutrons and electrons in a calcium ion, ⁴⁰Ca²⁺ (atomic number 20).

protons   = atomic number          = 20
neutrons  = mass number – protons  = 40 – 20 = 20
electrons = protons – charge       = 20 – 2  = 18

A 2+ ion has lost two electrons; the nucleus is unchanged.

Worked example 2. Gallium has two isotopes: 60% gallium-69 and 40% gallium-71. Calculate its relative atomic mass.

Ar = (60 × 69 + 40 × 71) ÷ 100
   = (4140 + 2840) ÷ 100
   = 69.8

The answer lies nearer the more common isotope, as a weighted average should.

2 and 3. Periodic relationships and atomic structure

Elements in the periodic table are arranged in order of atomic (proton) number. Elements with similar properties sit in columns called groups. It is called periodic because similar properties occur at regular intervals.

The reason is electronic structure. Electrons occupy the lowest available energy levels (shells): sodium is 2,8,1. Elements in the same group have the same number of outer electrons, which gives them similar chemical properties.

Group	Outer electrons	Behaviour	Trend down the group
1 (alkali metals)	1	React with oxygen, chlorine and water; form 1+ ions	Reactivity increases
7 (halogens)	7	Non-metals, diatomic molecules; form 1– ions	Reactivity decreases; melting and boiling points increase
0 (noble gases)	8 (helium 2)	Unreactive: stable arrangement of electrons	Boiling points increase

Metals are elements that react to form positive ions; they sit to the left and towards the bottom of the table. Non-metals do not form positive ions and sit towards the right and top.

Worked example 3. Explain why chlorine is more reactive than bromine.

1. Both have seven outer electrons, so both react by gaining one electron to form a 1– ion.
2. Chlorine atoms have fewer shells, so the outer shell is closer to the nucleus.
3. An incoming electron is more strongly attracted to the nucleus, so chlorine gains it more easily.

The same reasoning, reversed, explains Group 1: going down, the single outer electron is further from the nucleus, so it is lost more easily and reactivity increases.

4. Transferring or sharing electrons

There are three types of strong chemical bond. You explain each one in terms of electrostatic forces and the transfer or sharing of electrons.

• Ionic (metal + non-metal): outer electrons transfer from the metal atom to the non-metal atom. Ions from Groups 1, 2, 6 and 7 have the electronic structure of a noble gas. Oppositely charged ions attract in all directions in a giant lattice.
• Covalent (non-metals): atoms share pairs of electrons. The bonds are strong.
• Metallic (metals and alloys): outer electrons are delocalised and shared through a giant structure of positive ions.

Worked example 4. Describe the formation of magnesium oxide, and write the formula of potassium sulfide.

Mg  2,8,2  → loses 2 electrons → Mg²⁺  2,8
O   2,6    → gains 2 electrons → O²⁻   2,8
Formula: charges +2 and –2 balance → MgO

K is in Group 1 → K⁺ ;  S is in Group 6 → S²⁻
Two K⁺ balance one S²⁻ → K₂S

5. Shapes of molecules and giant structures

Bonding and structure together decide how a substance behaves. The stronger the forces between the particles, the higher the melting and boiling points.

Structure	Example	What is overcome on melting	Properties
Small molecules	methane, water	Weak intermolecular forces (not the covalent bonds)	Low melting and boiling points; no conduction
Polymers	poly(ethene)	Relatively strong intermolecular forces	Solids at room temperature
Giant covalent	diamond, graphite, silicon dioxide	Many strong covalent bonds	Very high melting points
Giant ionic	sodium chloride	Many strong electrostatic forces	High melting points; conduct when melted or dissolved
Metallic	copper	Strong metallic bonds	High melting points; conduct heat and electricity; malleable

Arrangement matters within one element too. In diamond each carbon forms four covalent bonds, so it is very hard and does not conduct. In graphite each carbon forms three bonds in hexagonal layers with no covalent bonds between layers, and one electron per carbon is delocalised, so graphite conducts. Alloys are harder than pure metals because different-sized atoms distort the layers, so they cannot slide as easily.

(Higher tier only) The particle model of solid spheres has limits: it shows no forces, and real particles are not solid spheres.

6. Barriers to reaction

Reactions can occur only when particles collide with sufficient energy. The minimum energy needed is the activation energy – the barrier on a reaction profile. Because the barrier differs between reactions, reactions go at different rates.

• Higher concentration, pressure or surface area → more frequent collisions → faster rate.
• Higher temperature → more frequent and more energetic collisions → faster rate.
• A catalyst provides a different pathway with a lower activation energy. It is not used up and does not appear in the equation.

Worked example 5. A reaction produces 48 cm³ of gas in 40 s. Calculate the mean rate.

mean rate = quantity of product formed ÷ time taken
          = 48 ÷ 40 = 1.2 cm³/s

7. Three ways reactions happen

Type	What moves	Examples in 8462
Proton transfer	H⁺ ions (a hydrogen ion is a proton)	Acids produce H⁺ in water; neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l)
Electron transfer	Electrons, from one species to another	Metals forming ions; displacement; electrolysis; metals with acids
Electron sharing	Pairs of electrons are shared in new bonds	Forming covalent molecules such as hydrogen chloride; halogens adding across C=C in alkenes

(Higher tier only) Oxidation is loss of electrons; reduction is gain of electrons. In Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, bromide ions lose electrons (oxidised) and chlorine gains them (reduced). At the cathode in electrolysis, positive ions gain electrons (reduction), for example 2H⁺ + 2e⁻ → H₂.

(Higher tier only) A strong acid is completely ionised in water; a weak acid is only partially ionised. As pH falls by one unit, the H⁺ concentration rises by a factor of 10. So a solution at pH 3 has 100 times the H⁺ concentration of one at pH 5.

8. Energy is conserved

Energy is neither created nor destroyed in a chemical reaction. The energy of the universe after the reaction equals the energy before. If a reaction transfers energy to the surroundings, the products must have less energy than the reactants by exactly that amount.

• Exothermic: transfers energy to the surroundings; temperature of the surroundings rises (combustion, neutralisation, hand warmers).
• Endothermic: takes in energy; temperature of the surroundings falls (thermal decomposition, citric acid with sodium hydrogencarbonate).

Worked example 6 (Higher tier only). Use these bond energies (kJ/mol) to calculate the energy change for 2H₂ + O₂ → 2H₂O: H–H 436, O=O 498, O–H 464.

bonds broken: 2 × 436 + 498 = 1370 kJ
bonds formed: 4 × 464       = 1856 kJ   (each H₂O has two O–H)
energy change = 1370 – 1856 = –486 kJ/mol  → exothermic

More energy is released forming bonds than is needed to break them.

Linking key ideas in one answer

Worked example 7. Explain why sodium chloride conducts electricity when molten but not when solid.

1. Sodium transfers one electron to chlorine, forming Na⁺ and Cl⁻ ions (idea 4).
2. The ions form a giant lattice held by strong electrostatic forces in all directions (idea 5).
3. In the solid the ions are fixed in place, so charge cannot flow.
4. When melted, the ions are free to move and carry charge.

Each step is a separate creditworthy point; join them with “so” or “because”.

Common errors

• Saying a 2+ ion has gained two protons – it has lost two electrons; the nucleus does not change.
• Explaining Group 1 reactivity with “more electrons” – the point is distance of the outer electron from the nucleus.
• Saying covalent bonds break when a simple molecular substance boils – the intermolecular forces are overcome.
• Writing “energy is made” or “energy is used up” – energy is transferred.
• Saying a catalyst “gives particles more energy” – it provides a pathway with a lower activation energy.
• (Higher tier only) Mixing up oxidation and reduction in terms of electrons: oxidation is loss.

Next steps

Test yourself with the practice questions, read AQA GCSE Chemistry exam preparation, and try the free 10-minute diagnostics.

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016, GCSE exams June 2018 onwards, version 1.1 (4 October 2019), published by AQA – section 4.11 Key ideas.