AQA GCSE Chemistry 8462: Key ideas – Revision Notes

These revision notes cover 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 key ideas are assessed across all papers, so they apply to Paper 1 and Paper 2 at Foundation and Higher tier, every May/June. Section 4.11 has no tier split, but content marked “Higher tier only” below is Higher tier only in the topics where it is taught.

For full explanations and worked examples, use the Key ideas study guide. Then test yourself with the Key ideas practice questions. The course hub is AQA GCSE Chemistry, and the printable checklist lists every topic. For a quick check of what you already know, try the free diagnostics.

The eight key ideas at a glance

#	Key idea	In one line	Taught in
1	Atoms and elements	Everything is atoms; about 100 naturally occurring elements	4.1.1
2	Periodic relationships	Similar properties recur at regular intervals, in groups	4.1.2
3	Explained by atomic structure	Same number of outer electrons → similar reactions	4.1.1.7, 4.1.2
4	Transfer or sharing	Ionic = transfer; covalent = shared pairs; metallic = delocalised	4.2.1
5	Shape and structure	Structure decides melting point, conduction, hardness	4.2.2, 4.2.3
6	Barriers to reaction	Activation energy → different rates	4.5.1.2, 4.6.1
7	Three ways to react	Proton transfer, electron transfer, electron sharing	4.4, 4.2.1.4
8	Energy conserved	Energy is transferred, never created or destroyed	4.5.1.1

Definitions to learn word for word

• Atom: the smallest part of an element that can exist.
• Element: a substance made of one type of atom; all its atoms have the same number of protons.
• Compound: two or more elements chemically combined in fixed proportions.
• Atomic number: number of protons. Mass number: protons + neutrons.
• Isotopes: atoms of the same element with different numbers of neutrons.
• Relative atomic mass: an average value that takes account of the abundance of the isotopes.
• Activation energy: the minimum amount of energy that particles must have to react.
• Catalyst: changes the rate of reaction but is not used up; gives a pathway with lower activation energy.
• Exothermic: transfers energy to the surroundings; surroundings get hotter.
• Endothermic: takes in energy from the surroundings; surroundings get colder.
• (Higher tier only) Oxidation is loss of electrons; reduction is gain of electrons.
• (Higher tier only) Strong acid: completely ionised in aqueous solution. Weak acid: only partially ionised.

Formulas and equations

What	Formula or equation	Tier
Neutrons	mass number – atomic number	Both
Electrons in an ion	protons – charge (for a 2– ion, add 2)	Both
Relative atomic mass	Σ(mass number × % abundance) ÷ 100	Both
Mean rate	quantity used or formed ÷ time taken (g/s or cm³/s)	Both
Neutralisation	H⁺(aq) + OH⁻(aq) → H₂O(l)	Both
Energy change	bonds broken – bonds formed (kJ/mol); negative = exothermic	Higher tier only
pH	1 unit lower → H⁺ concentration × 10	Higher tier only

Method in steps

Particles in an ion

1. Protons = atomic number (from the periodic table).
2. Neutrons = mass number – protons.
3. Electrons: subtract a positive charge, add a negative charge.

Formula of an ionic compound (Groups 1, 2, 6, 7)

1. Charge from group: Group 1 → +1, Group 2 → +2, Group 6 → –2, Group 7 → –1.
2. Use the smallest whole numbers of each ion so the charges add to zero.

Explaining a group trend

1. Same number of outer electrons → similar reactions.
2. Going down: more shells, outer electrons further from the nucleus.
3. Group 1: outer electron lost more easily → more reactive.
4. Group 7: an electron is gained less easily → less reactive.

Explaining a melting point

1. Name the structure (small molecules, giant covalent, giant ionic, metallic, polymer).
2. Name what must be overcome (intermolecular forces, covalent bonds, electrostatic forces, metallic bonds).
3. Say weak or strong, and many or few, then link to little or lots of energy.

Bond energy calculation (Higher tier only)

1. Draw or list every bond in the reactants and products.
2. Add the energy for bonds broken; add the energy for bonds formed.
3. Energy change = broken – formed. Keep the sign.

Structure and properties in one table

Structure	Melting point	Conducts?	Why
Small molecules	Low	No	Weak intermolecular forces; molecules have no overall charge
Polymers	Solid at room temperature	No	Intermolecular forces relatively strong between very large molecules
Giant covalent	Very high	No (graphite is the exception)	Many strong covalent bonds must be overcome
Giant ionic	High	Only when molten or dissolved	Ions must be free to move to carry charge
Metallic	Mostly high	Yes	Delocalised electrons carry charge and transfer thermal energy

Small worked reminders

• Formula from groups. Sodium (Group 1) and oxygen (Group 6): Na⁺ and O²⁻, so two Na⁺ for each O²⁻ → Na₂O.
• Reaction profile. Exothermic: products end lower than reactants. Endothermic: products end higher. The activation energy is the rise from the reactants to the top of the curve. A catalysed curve has a lower peak but the same start and end levels.
• Electron sharing. Hydrogen and chlorine form hydrogen chloride, H₂ + Cl₂ → 2HCl; each H–Cl bond is a shared pair.
• Electron transfer (Higher tier only). At the cathode, Cu²⁺ + 2e⁻ → Cu is a reduction because electrons are gained.
• Energy conserved. A hand warmer gets hot because the reaction transfers energy to the surroundings; the products store less energy than the reactants by the same amount.

Must-know distinctions

Pair	Tell them apart by
Ionic vs covalent	Transfer of electrons between metal and non-metal vs sharing of pairs between non-metals
Covalent bond vs intermolecular force	Strong bond inside a molecule vs weak force between molecules
Giant covalent vs small molecule	Many strong bonds to break (very high melting point) vs weak forces (low melting point)
Diamond vs graphite	Four bonds per carbon, no free electrons vs three bonds, one delocalised electron per carbon
Proton transfer vs electron transfer	H⁺ moves (acids, neutralisation) vs electrons move (redox, ions forming)
Exothermic vs endothermic	Temperature of surroundings rises vs falls
(Higher tier only) Strong vs concentrated	Degree of ionisation vs amount of acid per volume

Linking chains to learn

• Group → outer electrons → ion charge → formula. Magnesium is in Group 2, loses 2 electrons, forms Mg²⁺.
• Bonding → structure → property. Ionic → giant lattice → high melting point and conducts only when molten or dissolved.
• Collisions → activation energy → rate. Hotter → faster, more energetic particles → more collisions exceed the activation energy.
• Bonds broken and formed → energy transferred. Exothermic when forming bonds releases more energy than breaking them needs (Higher tier only).

Quick self-test

1. About how many different naturally occurring elements are there?
2. Boron is 20% boron-10 and 80% boron-11. Calculate its relative atomic mass.
3. Give the numbers of protons, neutrons and electrons in ³⁷Cl⁻ (atomic number 17).
4. How many electrons are in an Al³⁺ ion (atomic number 13)?
5. Write the formula of calcium chloride and calculate its Mr (Ca = 40, Cl = 35.5).
6. Why are the noble gases unreactive?
7. A reaction loses 3.6 g of gas in 90 s. Calculate the mean rate of reaction.
8. Name the three ways in which chemical reactions take place.
9. Which of the three ways describes neutralisation?
10. Why does graphite conduct electricity but diamond does not?
11. (Higher tier only) How many times greater is the H⁺ concentration at pH 2 than at pH 6?
12. (Higher tier only) In Zn + Cu²⁺ → Zn²⁺ + Cu, which species is oxidised?

Answers

1. About 100.
2. (20 × 10 + 80 × 11) ÷ 100 = 10.8
3. 17 protons, 20 neutrons, 18 electrons
4. 13 – 3 = 10 electrons
5. CaCl₂; Mr = 40 + 2 × 35.5 = 111
6. Their atoms have stable arrangements of electrons (full outer shells: eight, or two for helium).
7. 3.6 ÷ 90 = 0.04 g/s
8. Proton transfer, electron transfer, electron sharing.
9. Proton transfer: H⁺ ions from the acid react with OH⁻ ions to form water.
10. Graphite has one delocalised electron per carbon atom that can move and carry charge; in diamond all four outer electrons are used in covalent bonds.
11. 10⁽⁶⁻²⁾ = 10 000 times
12. Zinc (Zn loses two electrons to form Zn²⁺).

Where marks are usually lost

• Writing that an ion forms by gaining or losing protons instead of electrons.
• Giving “more shells” for Group 1 reactivity without saying the outer electron is further from the nucleus and less strongly attracted.
• Stating that boiling a simple molecular substance breaks covalent bonds.
• Forgetting that ionic compounds conduct only when molten or dissolved, not as solids.
• Saying higher temperature gives “more collisions” without also saying collisions are more energetic.
• Describing a catalyst as “lowering the energy of reactants” instead of providing a pathway with a lower activation energy.
• Writing that energy is “made” or “lost” in a reaction; it is transferred to or from the surroundings.
• (Higher tier only) Dropping the negative sign on an exothermic energy change, or counting only one O–H bond per water molecule.
• (Higher tier only) Confusing strong (fully ionised) with concentrated (lots of acid per volume).

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.