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Topics common to Paper 1 and Paper 2

Formulae, equations and hazards

0.1 - Recalling chemical formulaeLearn to recognise and write common element, compound and ion formulae. You will use the position of numbers, brackets and charges to distinguish formulae that contain similar symbols but represent different substances or particles.

0.2 - Word equationsTurn a description of a reaction into a word equation by separating the starting substances from the substances formed. You will practise keeping complete substance names together and including every reactant and product.

0.3 - Balancing chemical equationsWrite formula equations that conserve atoms, then add state symbols to show whether each substance is solid, liquid, gas or dissolved in water. Worked examples show how coefficients change particle numbers while the formulae stay fixed.

0.4 - Writing ionic equationsHigher tier: use a balanced formula equation to identify the particles that change and cancel unchanged spectator ions. You will check atoms and charge in precipitation and neutralisation equations; knowing common ion formulae and state symbols will help.

0.5 - Recognising chemical hazardsRecognise common pictograms on chemical containers and connect each warning to a suitable laboratory precaution. You will read the actual symbols and explain how a specific action reduces the danger they warn about.

0.6 - Assessing practical risks and precautionsAssess how a practical procedure could cause harm, then choose a precaution that reduces that particular risk. Worked examples link chemical and apparatus hazards to exposure, likelihood and seriousness across a range of practicals.

Topic 1 - Key concepts in chemistry

Atomic structure

1.1 - How atomic models changedDiscoveries of smaller particles changed the atom from Dalton’s indivisible sphere into a model with electrons surrounding a nucleus. You will connect each change to what scientists learned about atomic structure.

1.2 - Inside the atomLearn where protons, neutrons and electrons are found, and how a simple diagram represents their arrangement inside an atom.

1.3 - Charges and masses of subatomic particlesCompare the charges and masses of protons, neutrons and electrons. You will use the relative values correctly without confusing electric charge with mass.

1.4 - Why atoms are electrically neutralAn atom has no overall electrical charge. Learn why this requires equal numbers of protons and electrons, and why their masses and the number of neutrons do not set that balance.

1.5 - The sizes of atoms and nucleiThe nucleus is central but occupies only a tiny part of the atom. Learn to describe that size difference and recognise when a diagram enlarges the nucleus for clarity.

1.6 - Where an atom's mass is concentratedMost of an atom’s mass is concentrated in its tiny nucleus. Use particle masses to explain why occupying more space does not mean containing more mass.

1.7 - Understanding mass numberMass number counts the particles in an atom’s nucleus. Learn which particles to include, how to calculate the total, and what a name such as carbon-12 tells you.

1.8 - Atomic number and element identityAn element’s identity depends on its proton number. Learn to compare atoms using that number, and distinguish it from the total number of particles in the nucleus.

1.9 - IsotopesIsotopes let atoms belong to the same element while having different masses. Learn the two conditions for an isotope pair and use particle counts to identify one.

1.10 - Counting protons, neutrons and electronsUse atomic number and mass number to calculate protons, neutrons and electrons in neutral atoms. Then check a set of particle counts for mistakes.

1.11 - Why relative atomic masses are not always whole numbersA mass number counts particles in one nucleus, whereas relative atomic mass describes an average. Learn how isotope mixtures produce decimal averages without any atom containing a fraction of a neutron.

1.12 - Calculating relative atomic mass from isotopesFor this Higher-tier skill, calculate relative atomic mass as a weighted mean of isotope relative masses. You will use percentage and ratio abundances, then check and round the result.

The periodic table

1.13 - Mendeleev's periodic tableMendeleev used patterns in the properties of elements and their compounds to organise the elements known in his time. You will see how this evidence helped him decide which elements belonged together.

1.14 - Mendeleev's predictionsMendeleev used gaps in a repeating pattern to predict elements that had not yet been discovered. You will follow the reasoning from a gap to a prediction, then see how later discoveries tested it.

1.15 - Isotopes and the order of the periodic tableRelative atomic mass usually increases through the periodic table, but there are exceptions. You will use isotope abundance to explain why an element can have a larger average atomic mass than the element after it.

1.16 - Atomic number and position in the periodic tableAtomic number connects an element’s identity to its place in the periodic table. You will use the proton number of an atom to identify its element and compare positions in the table.

1.17 - Periods, groups and atomic numberThe periodic table uses atomic-number order, horizontal periods and vertical groups. You will use a table extract to recognise these patterns and identify elements likely to have similar chemical properties.

1.18 - Metals and non-metals in the periodic tableAn element’s position gives clues to whether it is a metal or a non-metal. You will connect the broad regions of the periodic table to outer-shell electrons, while recognising the limits of the pattern.

1.19 - Electron arrangements of the first 20 elementsAn electron arrangement shows how an atom’s electrons occupy shells. You will predict arrangements for neutral atoms of the first 20 elements, write them in numbers, and represent them in shell diagrams.

1.20 - Linking electron arrangements to the periodic tableFor the first 20 elements, a neutral atom’s electron arrangement connects to its period and group. You will use occupied shells and outer electrons to find a position, then work backwards from a position to an arrangement.

Ionic bonding

1.21 - Electron transfer and ionic bondingElectron transfer makes positive and negative ions. You will follow this process in sodium chloride and use a dot-and-cross model to distinguish making ions from the attraction that holds them together.

1.22 - Understanding ionsIons can be charged single atoms or charged groups of atoms. You will learn how to read their symbols and distinguish them from neutral atoms and molecules.

1.23 - Counting particles in ionsUse atomic number, mass number and ion charge to count protons, neutrons and electrons. Worked positive- and negative-ion examples show how the electron count changes while the nucleus stays the same.

1.24 - How ions form in ionic compoundsThe outer-shell electrons of Groups 1, 2, 6 and 7 help predict the ions formed in their ionic compounds. You will connect group number, electron transfer and charge, including cases where more than one atom supplies or receives electrons.

1.25 - Naming compounds: -ide and -ateThe endings -ide and -ate give clues about the elements in a compound. You will use familiar ionic compound names to interpret these clues and recognise why oxygen alone does not determine the ending.

1.26 - Writing formulae of ionic compoundsAn ionic compound formula records the simplest ratio of its ions. You will use supplied ion charges to balance positive and negative charge, then use brackets correctly when more than one group ion is needed.

1.27 - The ionic latticeIonic crystals contain a regular arrangement of positive and negative ions. You will connect this three-dimensional structure to the attractions between ions and interpret what a small, flat lattice diagram can show.

Covalent bonding

1.28 - Covalent bonds and shared electron pairsAtoms can join by sharing outer-shell electrons. This lesson explains what a shared pair means and how a diagram represents one covalent bond.

1.29 - How covalent bonds form moleculesCovalent bonds can join atoms into separate particles called molecules. You will distinguish atoms from molecules and recognise molecules of elements and compounds.

1.30 - The sizes of atoms and moleculesAtoms and small molecules are far too small to measure with a classroom ruler. You will learn their approximate scale, convert between metres and nanometres, and use that scale in a simple estimate.

1.31 - Drawing dot-and-cross diagramsA dot-and-cross diagram shows how outer-shell electrons are shared in a molecule. You will build and check diagrams for hydrogen, hydrogen chloride, water, methane, oxygen and carbon dioxide, then explain how sharing completes their outer shells.

Types of substance

1.32 - Classifying structures and bondingElements and compounds can be classified by the type of structure and bonding they have. This classification matters because structure and bonding explain physical properties such as relative melting point, boiling point, solubility in water and electrical conductivity. The best answers do not ju...

1.33 - Properties of ionic compoundsLearn why ionic compounds usually have high melting and boiling points, and why their electrical conductivity changes when they melt or dissolve. The key ideas are strong attractions and mobile ions.

1.34 - Properties of simple molecular substancesLearn how separate molecules explain low melting and boiling points and poor electrical conduction. Distinguish the strong bonds inside a molecule from the weaker forces between molecules.

1.35 - Carbon and giant covalent structuresLearn how graphite and diamond can be different substances while containing the same element, and what their giant covalent classification means.

1.36 - Structures of graphite and diamondCompare how carbon atoms join in diamond and graphite. Learn to read a structure model by counting neighbouring atoms and distinguishing a three-dimensional network from stacked sheets.

1.37 - How structure explains uses of diamond and graphiteUse the bonding in graphite and diamond to explain why graphite works as an electrode and lubricant, while diamond makes an effective cutting edge. Follow each structure → property → use link.

1.38 - Properties of fullerenes, C60 and grapheneCompare a hollow C₆₀ molecule with a one-atom-thick graphene sheet. Use the difference between separate molecules and a continuous network to explain their strength, electrical conduction and response to heating.

1.39 - Simple polymers and poly(ethene)Learn to recognise a simple polymer as very large molecules with carbon chains. Use poly(ethene) to read a short structural fragment and distinguish one bonded chain from a whole sample of polymer.

1.40 - Properties of metalsA copper wire carries current and a metal sheet can be pressed into shape. You will use the structure of a metal to explain both electrical conductivity and malleability, distinguishing electron movement from the sliding of layers.

1.41 - What scientific models can and cannot showChemical models make invisible structures easier to understand. You will compare dot-and-cross diagrams, ball-and-stick models and flat or three-dimensional representations, choosing what each can tell you and recognising what it leaves out.

1.42 - Comparing metals and non-metalsMetals and non-metals show useful patterns in their physical properties, but there are exceptions. You will describe the usual patterns and use several properties together to compare and classify elements.

Calculations involving masses

1.43 - Relative formula mass and percentage by massRelative formula mass is a way of adding up the relative atomic masses of all the atoms shown in a chemical formula. Once you know the relative formula mass, you can calculate what percentage of the compound's mass comes from one element.

1.44 - Empirical formulae from masses and percentagesMass data can be used to work out the formula of a simple compound. The important move is to convert masses, or percentages by mass, into a ratio of atoms. The formula you calculate in this way is an empirical formula: the simplest whole-number ratio of atoms of each element in the compound.

1.45 - Empirical and molecular formulaeEmpirical and molecular formulae both describe the elements in a compound, but they give different levels of detail. The key move is ratio: simplify a molecular formula to the smallest whole-number ratio, or scale an empirical formula up using relative molecular mass.

1.46 - Finding the formula of magnesium oxideAn empirical formula shows the simplest whole-number ratio of atoms of each element in a compound. In this experiment, magnesium is heated so that it reacts with oxygen to form magnesium oxide. The useful evidence is the mass of magnesium that reacted and the mass of oxygen that combined with it.

1.47 - Conservation of mass in closed and open systemsIn a chemical reaction, atoms are rearranged into new substances. The law of conservation of mass says that atoms are not made or destroyed in ordinary chemical reactions, so the total mass is conserved when every substance is included. The important GCSE skill is explaining why the measured mass...

1.48 - Masses from balanced equationsA balanced equation is not just a list of reactants and products. It gives the ratio in which substances react and form. In this Topic 1 calculation skill, you use that ratio with relative formula masses to calculate the mass of one reactant or product when the mass of another substance is given.

1.49 - Calculating mass concentrationA solution can contain a small or large mass of dissolved substance in the same volume. This lesson is about calculating concentration in g dm^-3, which means grams of solute per cubic decimetre of solution. You will convert volume units, calculate concentration and rearrange the relationship to...

1.50 - The mole and the Avogadro constantOne mole is a counting amount used in chemistry. This Higher-tier Topic 1 idea links a measurable mass in grams to a fixed number of particles. The key recall point is that one mole has two equivalent descriptions: it contains the Avogadro constant number of particles, and it has a mass equal to...

1.51 - Moles, particles and mass calculationsThis Higher-tier calculation lesson links the mass you can measure in the lab to the number of particles too small to count directly. The route is always built from two bridges: mass to moles, then moles to particles. Once those bridges are secure, the reverse calculations use the same relationsh...

1.52 - Limiting reactants and mass of productIn many reactions, one reactant is used up before the others. Once that reactant has gone, the reaction cannot keep making product, even if another reactant is still present. In this Higher-tier lesson, you will use mole ratios to find the limiting reactant and calculate the maximum mass of product.

1.53 - Finding reaction ratios from massesIn this Higher-tier lesson, you will turn measured masses into mole ratios and use them to deduce a balanced equation. You will learn how to handle fractional ratios and check that the equation fits both the data and conservation of atoms.

Topic 2 - States of matter and mixtures

States of matter

2.1 - Particle model of solids, liquids and gasesUse the particle model to compare the arrangement, movement and relative energy of solids, liquids and gases. These microscopic differences help explain why a solid keeps its shape, a liquid flows and a gas spreads out.

Methods of separating and purifying substances

Topic 3 - Chemical changes

Acids

Electrolytic processes

Topic 4 - Extracting metals and equilibria

Obtaining and using metals

Reversible reactions and equilibria

Topic 5 - Separate chemistry 1

Transition metals, alloys and corrosion

Quantitative analysis

The Haber process and fertilisers

Chemical cells and fuel cells

Topic 6 - Groups in the periodic table

Group 1 - Alkali metals

Group 7 - Halogens

Group 0 - Noble gases

Topic 7 - Rates of reaction and energy changes

Rates of reaction

Heat energy changes in chemical reactions

Topic 8 - Fuels and Earth science

Fuels

Earth and atmospheric science

Topic 9 - Separate chemistry 2

Chemical tests for gases and ions

Hydrocarbons

Polymers

Alcohols and carboxylic acids

Materials and nanoparticles

1.24 - How ions form in ionic compounds

1.24 - How ions form in ionic compounds

The outer-shell electrons of Groups 1, 2, 6 and 7 help predict the ions formed in their ionic compounds. You will connect group number, electron transfer and charge, including cases where more than one atom supplies or receives electrons.

Group number and ion charge

When a neutral atom loses electrons it becomes positive; when it gains electrons it becomes negative. Protons remain in the nucleus. For the Group 1 and 2 metals and Group 6 and 7 non-metals considered here, the usual simple ions have full outer electron shells.

Group	Outer-shell electrons	Change per atom	Usual ion charge	Example
1 metals	1	lose 1 electron	1+	$Na^+$
2 metals	2	lose 2 electrons	2+	$Mg^{2+}$
6 non-metals	6	gain 2 electrons	2−	$O^{2-}$
7 non-metals	7	gain 1 electron	1−	$Cl^-$

For example, sodium changes from 2,8,1 to 2,8, while chlorine changes from 2,8,7 to 2,8,8. This is a pattern for the specified groups in these compounds, not a rule for every element.

Accounting for every transferred electron

In these metal–non-metal compounds, electrons lost by the metal atoms are gained by the non-metal atoms. The resulting opposite charges attract electrostatically.

Magnesium oxide: one magnesium atom loses two electrons and forms $Mg^{2+}$ . One oxygen atom gains both electrons and forms $O^{2-}$ . The two electrons lost equal the two gained.

Calcium fluoride: one calcium atom loses two electrons and forms $Ca^{2+}$ . Each fluorine atom gains only one electron, so two fluorine atoms receive one electron each and form two $F^-$ ions. Do not turn one fluorine atom into $F^{2-}$ to make the numbers fit.

When two metal atoms are needed

Lithium oxide: each lithium atom loses one electron, while one oxygen atom gains two. Two lithium atoms therefore supply one electron each to one oxygen atom. The products are two $Li^+$ ions for each $O^{2-}$ ion.

Explain ion formation by naming the number and direction of transferred electrons and the charges produced. “The atoms want to be stable” does not describe these changes. A full outer shell helps describe the resulting electronic arrangement, but it does not replace the electron-transfer account.