How much of AQA A-Level Chemistry is section 3.2 Inorganic chemistry?

Inorganic chemistry is one of the three major sections (3.1 Physical, 3.2 Inorganic, 3.3 Organic) and is examined mainly in Paper 1, alongside relevant physical chemistry. The AS content (periodicity, Group 2, Group 7) and the A-level content (Period 3 oxides, transition metals, reactions of ions in aqueous solution) together carry a substantial share of marks, including descriptive trends, equations, observations and qualitative analysis.

What is the single best way to revise the trends in this topic?

Learn each trend as a cause, not just a direction. Almost every trend reduces to three factors: nuclear charge, atomic radius (distance) and shielding. Across a period nuclear charge rises and shielding is roughly constant; down a group radius and shielding rise. Use those three factors to explain atomic radius, ionisation energy, oxidising power of halogens and reducing power of halide ions, and you can reconstruct the answers under pressure.

Which equations must I be able to write from memory?

The Group 2 metals with water, the halogen displacement reactions, the sodium halides with concentrated sulfuric acid, the silver nitrate halide tests, the Period 3 elements with oxygen and the oxides with water, the ligand substitution reactions of copper, and the reactions of metal-aqua ions with hydroxide, ammonia and carbonate. Practise balancing them, including state symbols, because AQA marks these precisely.

Why are transition metal ions coloured?

When ligands bond to a transition metal ion, the 3d orbitals split into two energy levels. An electron absorbs visible light of the frequency that matches the energy gap and is promoted to the higher level; the complementary colour is transmitted, so the ion appears coloured. Changing the oxidation state, ligand or coordination number changes the size of the gap and so the colour. Ions with empty or full d sub-shells (Sc3+, Zn2+) are colourless.

How do the precipitate tests for metal ions actually work?

Metal ions in water are hexaaqua complexes that are slightly acidic. Adding a base (hydroxide or ammonia) removes hydrogen ions and precipitates the neutral hydroxide, whose colour identifies the metal (copper blue, iron(II) green, iron(III) brown, aluminium white). Adding the base in excess separates them further: aluminium hydroxide is amphoteric and redissolves in excess sodium hydroxide, while copper hydroxide redissolves in excess ammonia to a deep blue solution.

EnglandChemistry

AQA A-Level Chemistry 3.2 Inorganic chemistry: a deep dive on periodicity, the groups, Period 3, transition metals and aqueous ions

A deep-dive AQA A-Level Chemistry guide to section 3.2 Inorganic chemistry. Covers periodicity, Group 2, Group 7, the Period 3 elements and their oxides, transition metals, and reactions of ions in aqueous solution, with the trends, equations and exam patterns AQA repeats.

Generated by Claude Opus 4.824 min read3.2Updated 2026-06-02

Reviewed by: AI editorial process; not yet individually human-reviewed

Jump to a section

What section 3.2 actually demands
Periodicity (3.2.1)
Group 2, the alkaline earth metals (3.2.2)
Group 7, the halogens (3.2.3)
The Period 3 elements and their oxides (3.2.4)
Transition metals (3.2.5)
Reactions of ions in aqueous solution (3.2.6)
How section 3.2 is examined
Check your knowledge

What section 3.2 actually demands

Inorganic chemistry in AQA A-Level Chemistry is the study of the patterns in the periodic table: how and why properties change across a period and down a group, the characteristic reactions of the groups, and the rich behaviour of the transition metals and metal ions in water. The examiners test two linked skills: precise recall of trends, equations, colours and observations, and the ability to explain those trends from the underlying electron arrangement and bonding.

This guide walks through the six clusters of the topic in specification order, then sets out the exam patterns AQA repeats. Each cluster has a matching dot-point page with practice questions; this overview ties them together.

Periodicity (3.2.1)

Periodicity is the repeating pattern of properties across periods. Classify every element into the s, p, d or f block by the sub-shell holding its outermost electron. Across Period 3, atomic radius decreases because nuclear charge rises while shielding stays roughly constant, pulling the outer shell in. First ionisation energy generally increases for the same reason, but with two predictable dips: from magnesium to aluminium (an electron is removed from a higher-energy 3p orbital rather than 3s) and from phosphorus to sulfur (paired-electron repulsion in the 3p sub-shell of sulfur).

Melting point across a period reflects structure and bonding. Sodium, magnesium and aluminium are giant metallic, and the melting point rises as the ionic charge and number of delocalised electrons increase. Silicon is giant covalent and gives the peak. Phosphorus, sulfur and chlorine are simple molecular with weak van der Waals forces, and argon is monatomic, so all melt at low temperatures. Always name the structure when you explain a melting point.

Group 2, the alkaline earth metals (3.2.2)

Down Group 2, atomic radius increases, ionisation energy falls and reactivity increases, because the two outer electrons are lost more easily with more shielding and a larger radius. The metals react with water to form a hydroxide and hydrogen, more vigorously down the group; magnesium is slow with cold water but reacts well with steam to give the oxide.

The two solubility trends run in opposite directions, and this is heavily examined. Hydroxide solubility increases down the group, so magnesium hydroxide is only sparingly soluble while barium hydroxide is fairly soluble and strongly alkaline. Sulfate solubility decreases down the group, so barium sulfate is essentially insoluble, which underlies both the barium meal in medicine and the test for sulfate ions with acidified barium chloride. Learn the named uses: magnesium to extract titanium, calcium hydroxide to neutralise acidic soils, and Group 2 bases as antacids.

Group 7, the halogens (3.2.3)

Down Group 7, boiling point increases as the molecules get larger and the van der Waals forces strengthen, while electronegativity decreases. The halogens are oxidising agents whose power decreases down the group, because the larger atoms with more shielding gain an electron less easily. A more reactive halogen displaces a less reactive halide from solution, with colour changes you should know.

The halide ions are reducing agents whose power increases down the group, shown clearly by their reactions with concentrated sulfuric acid. Chloride gives only misty fumes of hydrogen chloride; bromide also reduces sulfur to sulfur dioxide and releases bromine; iodide reduces sulfur all the way to hydrogen sulfide and releases iodine. Halide ions are identified with acidified silver nitrate, giving white silver chloride, cream silver bromide and yellow silver iodide, distinguished further by their solubility in ammonia. Chlorine is used in water treatment, where it disproportionates with water to kill bacteria; the benefit outweighs the toxicity risk.

The Period 3 elements and their oxides (3.2.4)

Sodium reacts vigorously with cold water to give a strongly alkaline solution; magnesium is slow with cold water but reacts with steam. Burning the Period 3 elements in oxygen gives the oxides sodium oxide, magnesium oxide, aluminium oxide, silicon dioxide, phosphorus(V) oxide and the sulfur oxides.

The structure and bonding of the oxides change across the period: the metal oxides are giant ionic with high melting points, silicon dioxide is giant covalent with a very high melting point, and the non-metal oxides are simple molecular with low melting points. The ionic oxides dissolve to give alkaline solutions, the molecular oxides dissolve to give acids, and the pH of the resulting solutions falls across the period. Acid-base character mirrors this: the metal oxides are basic, aluminium oxide is amphoteric (reacting with both acids and bases), and the non-metal oxides are acidic.

Transition metals (3.2.5)

A transition metal is a d-block element that forms at least one stable ion with an incomplete d sub-shell, which excludes scandium and zinc. The partly filled 3d orbitals give the four signature properties: complex formation, coloured ions, variable oxidation states and catalytic activity.

A ligand is a species with a lone pair that forms a coordinate bond to the central ion; the coordination number is the number of such bonds, giving octahedral, tetrahedral, square planar or linear shapes. Complexes show cis-trans and optical isomerism (cisplatin is a key example). Ligands can be exchanged in substitution reactions, and multidentate ligands are favoured by the entropy-driven chelate effect. Colour arises from the splitting of the 3d orbitals: an electron absorbs visible light matching the energy gap and is promoted, and the complementary colour is transmitted; colorimetry uses this to find concentrations. Variable oxidation states arise from the similar 3d and 4s energies, and transition metals catalyse reactions either heterogeneously (adsorption on a surface, as in the Haber and Contact processes) or homogeneously (forming intermediates by cycling oxidation states).

Reactions of ions in aqueous solution (3.2.6)

Metal ions in water exist as hexaaqua complexes. The metal ion polarises the coordinated water, weakening the O-H bonds and releasing hydrogen ions, so the solution is acidic. The higher the charge density of the ion, the more acidic the solution, so 3+ ions are noticeably acidic while 2+ ions are only weakly so.

Adding a base removes hydrogen ions and precipitates the neutral hydroxide, whose colour identifies the metal: copper blue, iron(II) green, iron(III) brown and aluminium white. Adding the base in excess separates them: aluminium hydroxide is amphoteric and redissolves in excess sodium hydroxide as the aluminate ion, while copper hydroxide redissolves in excess ammonia to a deep blue complex. Carbonate ions distinguish 2+ from 3+ ions: 2+ ions give a carbonate precipitate, while acidic 3+ ions release carbon dioxide and precipitate the hydroxide. Together these reactions form the qualitative analysis scheme for identifying metal ions.

How section 3.2 is examined

A typical AQA profile for inorganic chemistry:

Trend explanations. Atomic radius, ionisation energy, oxidising and reducing power, and melting points, all explained from nuclear charge, radius and shielding, or from structure and bonding.
Equations and observations. Group 2 metals with water, halogen displacements, sodium halides with concentrated sulfuric acid, oxides with water, ligand substitution and aqua-ion reactions, with the colours and gases observed.
Qualitative analysis. Identifying halide and sulfate ions and the metal cations from precipitate colours and their behaviour in excess base.
Transition metal application. Complex shapes and isomerism, the origin of colour, colorimetry and catalysis.

Check your knowledge

A mix of recall and application questions covering the whole of section 3.2. Attempt them under timed conditions, then check against the solutions.

Explain why the first ionisation energy of sulfur is lower than that of phosphorus. (2 marks)
State and explain the trend in the solubility of the Group 2 sulfates. (2 marks)
Write an equation for the displacement reaction between chlorine and aqueous potassium iodide, and state the colour change. (3 marks)
Describe and explain what is observed when concentrated sulfuric acid is added to solid sodium bromide. (3 marks)
Explain why magnesium oxide has a much higher melting point than sulfur dioxide. (3 marks)
Explain, in terms of entropy, why EDTA readily replaces the six water ligands in a hexaaqua complex. (3 marks)
Explain why iron(III) solutions are more acidic than iron(II) solutions. (3 marks)
Describe how you would distinguish aluminium ions from copper(II) ions using sodium hydroxide and then ammonia. (4 marks)

Solutions

Q1: In phosphorus the 3p electron removed is from a half-filled sub-shell with all three electrons unpaired ( $3p^3$ ). In sulfur the 3p sub-shell has a paired electron ( $3p^4$ ), so there is repulsion between the two electrons in that orbital, making one easier to remove. Less energy is therefore needed, so sulfur has the lower first ionisation energy.
Q2: Solubility of the sulfates decreases down the group, so magnesium sulfate is soluble while barium sulfate is essentially insoluble. This is why barium sulfate can be used as a barium meal and why acidified barium chloride is the test for sulfate ions.
Q3: Chlorine is a stronger oxidising agent than iodine, so it displaces iodine: $Cl_2(aq) + 2KI(aq) \rightarrow 2KCl(aq) + I_2(aq)$ . The colourless solution turns brown (or a black solid of iodine may form).
Q4: Bromide is a moderately strong reducing agent. It first gives misty fumes of hydrogen bromide ( $HBr$ ), then reduces the sulfur in sulfuric acid from $+6$ to $+4$ , giving sulfur dioxide ( $SO_2$ , a choking gas) and orange-brown bromine ( $Br_2$ ). Observations: misty fumes, a choking gas and brown fumes.
Q5: Magnesium oxide is giant ionic, so melting it breaks many strong electrostatic forces between $Mg^{2+}$ and $O^{2-}$ ions, needing a lot of energy. Sulfur dioxide is simple molecular with only weak van der Waals forces between molecules, which need little energy to overcome. So $MgO$ has the much higher melting point.
Q6: The reaction $[M(H_2O)_6]^{n+} + EDTA^{4-} \rightarrow [M(EDTA)]^{(n-4)} + 6H_2O$ converts two species into seven, increasing the number of free particles. The entropy increases ( $\Delta S$ positive), while the enthalpy change is small because similar bonds break and form. A positive $\Delta S$ makes $\Delta G$ negative, so the substitution is feasible (the chelate effect).
Q7: $Fe^{3+}$ has a higher charge and smaller radius than $Fe^{2+}$ , so a greater charge density. It polarises the coordinated water molecules more strongly, weakening the O-H bonds and releasing more $H^+$ ions. The higher $H^+$ concentration makes the iron(III) solution more acidic.
Q8: With sodium hydroxide, both give a precipitate (white $Al(OH)_3$ , blue $Cu(OH)_2$ ); adding excess $NaOH$ redissolves only the amphoteric $Al(OH)_3$ to a colourless solution, while $Cu(OH)_2$ stays blue. With ammonia, both give a precipitate; adding excess ammonia redissolves $Cu(OH)_2$ to a deep blue solution, while $Al(OH)_3$ stays as a white precipitate. So $Al^{3+}$ dissolves in excess $NaOH$ and $Cu^{2+}$ dissolves in excess ammonia.

Sources & how we know this

AQA A-level Chemistry (7405) specification — AQA (2015)