CCh

Chemistry

AQA · Paper 1 · Topics C1–C5

Triple = separate chemistry only. HT = higher tier only. If you're on combined science, skip anything marked Triple.

C1At

Atomic structure & the periodic table

Particle	Relative mass	Relative charge	Where
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	Very small	−1	Shells

Atomic number (bottom) = protons. In a neutral atom, protons = electrons.
Mass number (top) = protons + neutrons. So neutrons = mass − atomic. ²³₁₁Na → 11 p, 11 e, 12 n.
Isotopes: same protons, different neutrons.
Ions: same protons, different electrons. Lose e⁻ → positive; gain e⁻ → negative. Mg²⁺ has 10 electrons.
Atoms have no overall charge because protons = electrons.

Relative atomic mass (Ar)

Ar = Σ(abundance × mass number) ÷ total abundance

Chlorine: 75% ³⁵Cl and 25% ³⁷Cl → (75×35 + 25×37) ÷ 100 = 35.5

History of the atom

Dalton: atoms are tiny solid spheres that can't be divided.
Thomson discovered the electron → plum pudding model (ball of positive charge with electrons embedded).
Rutherford, alpha scattering: most alpha particles went straight through → mostly empty space. A few deflected → positive charge concentrated in a tiny nucleus. Nuclear model replaced plum pudding.
Bohr: electrons orbit the nucleus at fixed distances (shells).
Protons found in the nucleus; Chadwick proved neutrons exist.

Discovery order: electron → proton → neutron.

Groups

Group 1 alkali metals: 1 outer electron. React with water → metal hydroxide + hydrogen. More reactive going down: outer electron further from nucleus, weaker attraction, lost more easily.
Group 7 halogens: 7 outer electrons, exist as pairs (Cl₂). Less reactive going down: harder to gain an electron. A more reactive halogen displaces a less reactive one: chlorine + potassium bromide → potassium chloride + bromine.
Group 0 noble gases: full outer shell, so unreactive. Boiling points increase going down.

Watch out: some books call it "Group 8". Write Group 0 in the exam.

Separating mixtures

Method	Use it to…	Key details
Filtration	Remove an insoluble solid from a liquid	Filtrate passes through; residue stays on the paper
Evaporation	Get a soluble solid from a solution	Quick, but some solids may decompose
Crystallisation	Get pure crystals of a soluble salt	Heat gently until crystals start to form, cool, filter, dry
Simple distillation	Get the solvent from a solution	Solvent boils, vapour cools in condenser, liquid collected
Fractional distillation	Separate liquids with different boiling points	Fractionating column; lowest boiling point collected first
Chromatography	Separate/identify dyes	Pencil baseline above solvent; Rf = spot ÷ solvent distance

Solute dissolves in solvent to make a solution.

C2Bd

Bonding & structure

Finding bonding hard? Open the Bonding deep dive at the end of this section. It teaches C2 step by step with diagrams.

Bond	Between	What happens
Ionic	Metal + non-metal	Electrons transferred → oppositely charged ions attract
Covalent	Non-metal + non-metal	Pairs of electrons shared
Metallic	Metal atoms	Positive ions in a sea of delocalised electrons; strong attraction between them

Making NaCl: sodium loses 1 electron → Na⁺; chlorine gains 1 electron → Cl⁻. Both now have a full outer shell. The ions are held together by strong electrostatic attraction.

Giant ionic lattice (e.g. NaCl)

Strong electrostatic forces between oppositely charged ions, in all directions.
High melting/boiling points: lots of energy needed to overcome them.
Conduct only when molten or dissolved: ions are free to move and carry charge.

Simple molecular (H₂O, CO₂, CH₄, O₂)

Low melting/boiling points: only weak intermolecular forces are overcome, not the covalent bonds.
Don't conduct: no free electrons or ions, no overall charge.

Giant covalent (diamond, graphite, silicon dioxide)

Very high melting points: many strong covalent bonds must break.
Diamond: each C has 4 covalent bonds → very hard, doesn't conduct.
Graphite: each C has 3 bonds, in layers. Weak forces between layers → slide, soft. One delocalised electron per C → conducts.

Polymers

Many small monomers join into long-chain polymers.

Metals & alloys

High melting points, good conductors (delocalised electrons).
Malleable: layers of atoms slide over each other.
Alloys are harder: different-sized atoms distort the layers, so they can't slide as easily.

Bonding deep dive: step by step

The 5 sentences that score most C2 marks

Learn these word for word. Almost every C2 "explain" question is one of these.

Giant ionic: strong electrostatic forces between oppositely charged ions need a lot of energy to overcome.→ high melting point
Simple molecular: weak intermolecular forces between molecules need little energy to overcome.→ low melting/boiling point. The covalent bonds do NOT break.
Giant covalent: many strong covalent bonds need a lot of energy to break.→ very high melting point (diamond, graphite, silicon dioxide)
Conducting: the ions / delocalised electrons are free to move and carry charge.→ ions for ionic compounds, electrons for metals and graphite
Alloys: different-sized atoms distort the layers, so they can't slide over each other.→ alloys are harder than pure metals

Which bonding? Ask what it's made of

Metal + non-metal
NaCl, MgO, CaCl₂→Ionic
giant ionic lattice

Non-metals only
H₂O, CO₂, CH₄, Cl₂→Covalent
usually simple molecules

The 4 exceptions
diamond, graphite, graphene, silicon dioxide→Covalent
giant covalent

Metal only, or alloy
Cu, Fe, steel→Metallic
ions in a sea of electrons

Shortcut: if it's covalent and not on the list of 4 giant ones, assume it's a simple molecule.

How the bonds form: dot-and-cross

Ionic: electrons are transferred

Sodium (2,8,1) loses its 1 outer electron → Na⁺ (2,8).
Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8). The cross is the electron from sodium.
Both now have a full outer shell. Draw ions in square brackets with the charge outside.
Magnesium oxide: Mg (2,8,2) loses 2 electrons → Mg²⁺; O (2,6) gains 2 → O²⁻.

Working out formulas: make the charges cancel. Mg²⁺ and Cl⁻ need two Cl⁻ → MgCl₂. Al³⁺ and O²⁻: 2 × 3+ = 6+ and 3 × 2− = 6− → Al₂O₃.

Covalent: electrons are shared

Each covalent bond = one shared pair of electrons, drawn where the circles overlap.
Water: O shares one pair with each H → 2 bonds. Methane CH₄ → 4 bonds. Cl₂ → 1 bond.
Only draw the outer shells.

Limits of these diagrams (a common 1-mark question): dot-and-cross diagrams don't show the 3D arrangement or the relative sizes of atoms.

The four structures

Giant ionic lattice

Picture: Huge 3D grid of + and − ions, alternating.
Melting: High. Strong electrostatic forces between oppositely charged ions, in all directions.
Conducts: Only molten or dissolved, when the ions are free to move. Never as a solid.
e.g.: Sodium chloride, magnesium oxide

Simple molecules

Picture: Separate small molecules. Strong bonds inside, weak forces between.
Melting: Low. Only the weak intermolecular forces are overcome.
Conducts: Never. No overall charge, no free electrons or ions.
e.g.: Water, CO₂, methane, oxygen, chlorine

Biggest C2 mark-loser: "the covalent bonds break when it boils". They don't. Only the forces between molecules are overcome. Bigger molecules have stronger intermolecular forces, so higher boiling points.

Giant covalent: diamond vs graphite

Both: Very high melting point: many strong covalent bonds to break.
Diamond: Each C has 4 bonds → rigid, very hard, doesn't conduct (no free electrons). Used in cutting tools.
Graphite: Each C has 3 bonds, in hexagon layers. Weak forces between layers → slide → soft, slippery. 1 delocalised electron per C → conducts.
Graphene: A single layer of graphite. Strong, conducts.
Fullerenes: Hollow carbon molecules (C₆₀ ball, nanotubes). Nanotubes are strong for their size; used in electronics.

Metallic

Picture: Layers of positive metal ions in a sea of delocalised electrons.
Melting: High. Strong attraction between positive ions and delocalised electrons.
Conducts: Yes, solid and liquid. Delocalised electrons move and carry charge (and heat).
Shape: Malleable: layers slide. Alloys are harder:

Master table: screenshot this

Structure	Melting point	Conducts solid?	Conducts molten / dissolved?
Giant ionic	High	No	Yes (ions move)
Simple molecular	Low	No	No
Giant covalent	Very high	No (graphite yes)	–
Metallic	High	Yes (electrons)	Yes

Other C2 bits that come up

Polymers: very large molecules. Stronger intermolecular forces than small molecules → solid at room temperature.
States of matter: melting and boiling happen at the same temperature as freezing and condensing. More energy needed = stronger forces between particles.
State symbols: (s) solid, (l) liquid, (g) gas, (aq) dissolved in water.

C3Mo

Quantitative chemistry

Ar = relative atomic mass (big number on the periodic table).
Mr = sum of the Ar of every atom in the formula.

MgCl₂ = 24 + (2 × 35.5) = 95 Ca(OH)₂ = 40 + 2×(16+1) = 74

Moles

moles = mass ÷ Mr | mass = moles × Mr

1 mole = 6.02 × 10²³ particles (Avogadro's constant).

Mass from an equation: 4 steps

Moles of what you know = mass ÷ Mr
Use the big numbers for the ratio
Moles of what you want
Mass = moles × Mr

2NaClO₃ → 2NaCl + 3O₂ : 0.4 mol NaClO₃ → 0.6 mol O₂ → 0.6 × 32 = 19.2 g

Limiting reactant: used up first, so it limits how much product forms. The other is in excess.

Conservation of mass

Mass seems to increase if a reactant is a gas from the air (Mg + O₂ from air → MgO).
Mass seems to decrease if a product is a gas that escapes (CaCO₃ → CaO + CO₂).

Concentration

concentration (g/dm³) = mass (g) ÷ volume of solution (dm³)

Watch out: it's volume of solution, not solvent. cm³ → dm³: ÷ 1000.

C4pH

Chemical changes

0 strong acid7 neutral14 strong alkali

Universal indicator: red/orange acid, green neutral, blue/purple alkali.
Acids → H⁺ ions in water. Alkalis → OH⁻ ions in water.
Base: neutralises an acid → salt + water. Alkali: a base that dissolves in water.
Neutralisation: H⁺ + OH⁻ → H₂O
Strong acid: completely ionises in water (HCl, H₂SO₄, HNO₃). Weak: partially ionises (ethanoic acid).

Acid reactions and naming salts

acid + metal → salt + hydrogen
acid + metal oxide/hydroxide → salt + water
acid + metal carbonate → salt + water + carbon dioxide
Hydrochloric → chloride; sulfuric → sulfate; nitric → nitrate. Zinc + hydrochloric acid → zinc chloride + hydrogen.

Making a soluble salt (required practical)

Copper oxide + sulfuric acid → copper sulfate + water. Warm acid → add oxide until excess → filter → heat filtrate to evaporate some water → leave to crystallise → dry.

Redox: OIL RIG

Oxidation Is Loss of electrons. Reduction Is Gain.
Mg + CuSO₄ → MgSO₄ + Cu: Mg loses electrons (oxidised), Cu²⁺ gains them (reduced).

Electrolysis

Electrolyte must be molten or dissolved so ions can move.
Positive ions → cathode (−), gain electrons → reduced.
Negative ions → anode (+), lose electrons → oxidised.
Molten lead bromide: Pb²⁺ + 2e⁻ → Pb at cathode; 2Br⁻ → Br₂ + 2e⁻ at anode.

Trick: PANIC – Positive Anode, Negative Is Cathode.

Reactivity series and extracting metals

K · Na · Li · Ca · Mg · C · Zn · Fe · H · Cu · Ag · Au

More reactive metals lose electrons (form positive ions) more easily. They react faster with water and acid.
A more reactive metal displaces a less reactive one from its compound: magnesium + copper sulfate → magnesium sulfate + copper.
Metals below carbon (zinc, iron, copper) are extracted by heating the oxide with carbon. The oxide loses oxygen, so it's reduced.
Metals above carbon (e.g. aluminium) are extracted by electrolysis, which uses a lot of energy.
Gold is so unreactive it's found as the metal itself.
Oxidation can also mean gaining oxygen; reduction can mean losing oxygen.

Electrolysis of solutions

Water gives extra H⁺ and OH⁻ ions.
Cathode: the metal forms only if it's less reactive than hydrogen; otherwise hydrogen.
Anode: a halogen forms if there are halide ions; otherwise oxygen.
See the Practicals tab for the full method and gas tests.

C5En

Energy changes

Exothermic: energy released to surroundings, temperature rises. Burning, neutralisation, hand warmers.
Endothermic: energy taken in, temperature falls. Thermal decomposition, sports ice packs.
Breaking bonds takes in energy. Making bonds releases energy.
Activation energy: minimum energy needed for particles to react.

Energy change = bonds broken − bonds made (negative → exothermic)

H₂ + Cl₂ → 2HCl (H–H 436, Cl–Cl 243, H–Cl 432): broken = 679, made = 2 × 432 = 864 → 679 − 864 = −185 kJ/mol

Reaction profiles

Exothermic: products lower than reactants. Endothermic: products higher.
Activation energy: reactants up to the peak. Overall change: reactants to products.

Temperature change practical

Polystyrene cup with a lid (insulates, less heat loss), thermometer, measure start and highest/lowest temperature.

Required practicals are worth around 15% of the marks, and 6-mark "describe a method" questions usually come from them. For each one, know the method, the variables and the safety points.

RP · Making a soluble salt

Example: copper sulfate from copper oxide and sulfuric acid. CuO + H₂SO₄ → CuSO₄ + H₂O

Measure 40 cm³ of dilute sulfuric acid into a beaker and warm it gently with a Bunsen burner.
Add copper oxide one spatula at a time, stirring, until no more dissolves. The oxide is now in excess, so all the acid has reacted.
Filter to remove the excess copper oxide. Collect the blue filtrate in an evaporating basin.
Heat the filtrate on a water bath until crystals start to form at the edge, then stop.
Leave to cool and crystallise. Filter off the crystals and pat them dry with filter paper.

Why excess?: So all the acid is used up and the salt isn't contaminated with acid.
Why warm?: Speeds up the reaction.
Safety: Eye protection (acid irritant, copper sulfate harmful); don't heat to dryness (crystals can spit).

RP · Titration Triple

Use a pipette and filler to measure exactly 25.0 cm³ of alkali into a conical flask.
Add a few drops of indicator (phenolphthalein or methyl orange). Stand the flask on a white tile.
Fill a burette with acid and record the starting volume at eye level.
Add acid, swirling, then drop by drop near the end, until the indicator just changes colour (the end point).
Record the final volume. Titre = final − start.
Repeat until you get concordant results (within 0.10 cm³ of each other) and calculate a mean.

moles = concentration (mol/dm³) × volume (dm³) HT

Indicator: Phenolphthalein: pink in alkali → colourless. Methyl orange: yellow in alkali → red in acid.
White tile: To see the colour change clearly.

Don't use universal indicator: it changes colour gradually, so there's no sharp end point.

RP · Electrolysis of aqueous solutions

Pour the solution (e.g. copper chloride, sodium chloride) into a beaker.
Put in two inert graphite electrodes that don't touch, connected to a low-voltage d.c. supply.
Collect any gas in inverted test tubes over the electrodes. Observe each electrode.
Test the gases.

What forms where

Electrode	Rule
Cathode (−)	The metal forms if it's less reactive than hydrogen (e.g. copper). Otherwise hydrogen forms.
Anode (+)	A halogen forms if the solution has halide ions (Cl⁻, Br⁻, I⁻). Otherwise oxygen forms.

Gas tests

Gas	Test	Result
Hydrogen	Lighted splint	Squeaky pop
Oxygen	Glowing splint	Relights
Chlorine	Damp blue litmus paper	Bleached white (may turn red first)
Carbon dioxide	Bubble through limewater	Turns cloudy

Copper chloride solution: copper at the cathode, chlorine at the anode. Sodium chloride solution: hydrogen at the cathode, chlorine at the anode.

RP · Temperature changes

Measure 30 cm³ of hydrochloric acid into a polystyrene cup, standing in a beaker so it doesn't tip.
Record the starting temperature with a thermometer.
Add 5 cm³ of sodium hydroxide solution, stir, put the lid on and record the highest temperature.
Repeat with 10, 15, 20… cm³ of sodium hydroxide.
Repeat the whole experiment and calculate means. Plot temperature rise against volume.

IV: Volume of sodium hydroxide
DV: Temperature change
Control: Volume and concentration of acid, same cup, starting temperature

Polystyrene and a lid reduce energy lost to the surroundings.
The temperature rise peaks when the acid is exactly neutralised, then falls as extra cold alkali is added.

Practical words you must know

Word	Meaning
Independent variable	The one you change
Dependent variable	The one you measure
Control variable	Kept the same so it's a fair test
Repeatable	Same person, same method → similar results
Reproducible	Different person or method → similar results
Accurate	Close to the true value
Precise	Repeat results close together
Concordant	Titres within 0.10 cm³ of each other
Anomalous	Doesn't fit the pattern. Leave it out of the mean.

These are the question types that come up on every paper. Learn the pattern and you'll pick up marks even on topics you're shaky on.

Unfamiliar context? Don't panic

Lots of marks are on things you haven't been taught directly: an element you've never met, a new compound, a new reaction. They're always testing a topic you do know.

Find the topic: an element at the bottom of Group 1 = group trends. A substance with a weird melting point = structure and bonding. A new equation = moles and ratios.
Use the pattern: if it's in the same group, it behaves like the others, just more or less reactive.
Use the data given. Quote numbers from the table. The examiner put them there for you.
"Predict" and "suggest" mean any sensible answer based on a pattern can score. Always write something.

Command words: what they want

Word	Do this
Give / State / Name	A short answer. No explanation.
Describe	Say what happens or what it's like. No "why" needed.
Explain	Say why. Link each point with "because" or "so".
Compare	Similarities and/or differences. Use "whereas".
Calculate	Show working; give units.
Suggest	Apply what you know to something new. Any sensible answer can score.

Marks = points. A 3-mark question needs 3 separate correct ideas.

"Explain why X has a high/low melting point"

Name the structure (giant ionic, simple molecular, giant covalent, metallic).
Name what holds it together (electrostatic forces between ions / covalent bonds / intermolecular forces).
Say whether they're strong or weak.
Link to energy: a lot / little energy needed to overcome them.

Example
Methane is a simple molecule with weak intermolecular forces between molecules, so little energy is needed to overcome them.

Mark killer: saying covalent bonds break when a simple molecule melts or boils. That gets 0.

"Explain why X conducts (or doesn't)"

Name the charged particle: ions or delocalised electrons.
Say they're free to move.
Say they carry charge (current) through the substance.

Example
Molten NaCl conducts because the ions are free to move and carry charge. Solid NaCl doesn't because the ions are fixed in the lattice.

Mark killer: "electrons move" for an ionic compound. In ionic compounds it's always the ions.

Group 1 and Group 7 reactivity

Group 1, more reactive going down:

More shells, so outer electron is further from the nucleus.
Weaker attraction to the nucleus (more shielding).
Outer electron is lost more easily.

Group 7, less reactive going down: same first two points, but they need to gain an electron, so it gets harder.

Mark killer: "the atom is bigger so it's more reactive". You need the attraction and losing/gaining an electron.

Mole calculations from an equation

Write the Mr of what you know and what you want.
Moles of known = mass ÷ Mr.
Use the balancing numbers for the ratio.
Mass of unknown = moles × Mr.

Example
4.8 g Mg in Mg + 2HCl → MgCl₂ + H₂. Moles Mg = 4.8 ÷ 24 = 0.2. Ratio 1:1, so 0.2 mol MgCl₂. Mass = 0.2 × 95 = 19 g.

Always show working. Even with the wrong final answer you can get 2–3 marks out of 4.

Concentration and unit traps

Convert cm³ to dm³: ÷ 1000.
Concentration = mass ÷ volume.
Rearrange when needed: mass = concentration × volume.

Example
5.0 g in 200 cm³ → 0.200 dm³ → 5.0 ÷ 0.200 = 25 g/dm³.

Bond energy calculations

Draw out or list every bond in the reactants. Add them up (bonds broken).
Do the same for the products (bonds made). Remember the big numbers: 2HCl = 2 H–Cl bonds.
Energy change = broken − made.
Negative means exothermic, positive means endothermic.

6-mark method questions

Marked in levels, not points. To get Level 3 (5–6 marks) you need a complete method in a logical order that someone else could follow.

Write numbered steps, one action per step.
Name equipment (beaker, filter funnel, evaporating basin, Bunsen).
Say why for the key steps ("add copper oxide until in excess so all the acid reacts").
Finish the job: get to dry crystals or a pure product.

Learn the two classics: salt from rock salt and copper sulfate from copper oxide. Both are in the paper.

Electrolysis questions

Metal (or hydrogen) goes to the cathode (−). Non-metal to the anode (+).
Cathode: ions gain electrons = reduction. Anode: ions lose electrons = oxidation.
Half equations: count charge. Pb²⁺ needs 2e⁻. Two Br⁻ make Br₂ and give 2e⁻.
Must be molten or dissolved so ions are free to move.