AQA A-Level Chemistry: Inorganic Chemistry — Complete Revision Guide (7405)

Inorganic chemistry on AQA 7405 is the section where the ideas you built up in atomic structure, bonding and energetics finally cash out as predictions about real elements. The whole arc — from periodicity across Period 3, through the chemistry of Groups 2 and 7, into the colourful world of the transition metals — is one extended exercise in using electron configurations, ionic and covalent character, redox and complex-ion equilibria to explain what each element actually does in the lab. It is also the section that examiners reward most for precise, structured language: "the ionic radius decreases" beats "it gets smaller" every time.

This guide walks through the ten inorganic sub-topics on the AQA specification. For each you will find the core ideas, the underlying physical reasoning, the common pitfalls, and a link into the matching lesson in the LearningBro AQA A-Level Chemistry Inorganic course.

Synoptic Preview: How Inorganic Connects to the Rest of 7405

Inorganic chemistry is the most synoptic block on AQA 7405. Almost every quantitative argument here borrows machinery from elsewhere in the specification.

• Bonding and structure (bonding course) explains why Period 3 melting points peak at silicon: lattice and covalent structures, not atomic mass, set the trend. The same lattice-energy reasoning underpins the relative thermal stabilities of Group 2 carbonates and nitrates.
• Energetics (energetics course) supplies the Born-Haber cycle that quantifies the lattice enthalpies of the Group 2 halides and oxides, and the enthalpy of hydration that drives Group 2 solubility trends.
• Redox and electrochemistry (redox course) provides the E° framework used for transition-metal redox couples — Fe(II)/Fe(III), Mn(VII)/Mn(II), Cr(VI)/Cr(III) — and explains why a given oxidising agent oxidises some halides but not others.
• Acids, bases and buffers (physical course) explains why hexaaqua ions of M(III) metals are markedly acidic: the high charge density of the central ion polarises coordinated water, weakening the O-H bond. The same Ka framework predicts which carbonate-precipitate or hydroxide-precipitate reactions occur.
• Analytical chemistry (analytical course) closes the loop: every "identify the unknown" question chains aqueous-ion tests with the precipitates and colours you meet in this section.

Time spent on inorganic is therefore rewarded twice — once on Paper 1 directly, and again on every synoptic question that asks why a particular trend exists.

Guide Overview: Ten Inorganic Sub-Topics

The ten lessons in this guide map directly onto the AQA inorganic content:

1. Periodicity and trends across Period 3
2. Group 2 — the alkaline earth metals
3. Group 7 — the halogens
4. Period 3 elements and their oxides
5. Transition metals — general properties
6. Complex ions and ligands
7. Ligand substitution reactions
8. Colour in transition-metal complexes
9. Reactions of aqueous transition-metal ions
10. Transition-metal catalysis

Three of the AQA Required Practicals — RP10 (Group 2 and Group 7 trends), RP11 (preparation of a transition-metal complex and tests on aqueous ions) and RP12 (ligand-substitution colour changes) — sit inside this block, so practical-style questions on Paper 3 lean heavily on inorganic content.

What the AQA 7405 Specification Covers

AQA A-Level Chemistry (7405) is assessed through three papers: Paper 1 (Inorganic and Physical, 2h, 105 marks), Paper 2 (Organic and Physical, 2h, 105 marks) and Paper 3 (all content + Required Practicals, 2h, 90 marks). Inorganic is examined heavily on Paper 1 and reappears on Paper 3 through RP10-RP12 contexts.

Sub-topic	Spec area	Typical Paper 1 weight
Periodicity across Period 3	§3.2.1	3-5 marks
Group 2 alkaline earth metals	§3.2.2	4-6 marks
Group 7 halogens	§3.2.3	5-8 marks
Period 3 elements and oxides	§3.2.4	4-6 marks
Transition metals general	§3.2.5	3-5 marks
Complex ions and ligands	§3.2.5	4-6 marks
Ligand substitution	§3.2.5	4-6 marks
Colour in complexes	§3.2.5	3-5 marks
Aqueous-ion reactions	§3.2.6	6-10 marks
Transition-metal catalysis	§3.2.5	3-5 marks

These weights are estimates modelled on recent AQA paper formats. What is reliable is that a long aqueous-ion question — typically eight to twelve marks — appears on essentially every Paper 1, and that a colour-change identification table is a frequent Paper 3 anchor.

Periodicity and Trends Across Period 3

Periodicity is the AQA section that asks you to read trends straight off the periodic table. Across Period 3 (Na → Ar), atomic radius decreases because nuclear charge increases while electrons are added to the same principal shell, so the effective nuclear charge felt by the outer electrons grows.

First ionisation energy generally increases across the period, with two characteristic dips: between Mg and Al (the 3p electron in Al is at higher energy than the 3s electrons in Mg and easier to remove) and between P and S (in S the first 3p electron to leave comes from a doubly-occupied orbital, where electron-electron repulsion lowers the ionisation energy). These two anomalies are reliable mark-earners and were rationalised in modern form by Pauling's analysis of orbital energies.

Melting points trace the bonding type. Na, Mg and Al are metallic, with melting point rising as charge on the cation rises and the metallic lattice strengthens. Si is a giant covalent (macromolecular) lattice and has the highest melting point in the period. P4, S8 and Cl2 are simple molecular solids held by van der Waals forces, with S8 melting highest of the three because the larger molecule has more induced-dipole-induced-dipole interactions. Ar is a monatomic gas.

A common pitfall is to conflate "bond strength" with "intermolecular force strength" for molecular elements — phosphorus melts because P4 molecules separate, not because P-P bonds break. See the periodicity and trends lesson.

Group 2: Alkaline Earth Metals (RP10 Anchor)

The Group 2 metals (Be, Mg, Ca, Sr, Ba) are reducing agents whose chemistry is dominated by the ease of losing two outer s electrons to form M(2+) ions. Down the group, atomic radius increases, first and second ionisation energies fall, and reducing power rises — so reactivity with water increases from Mg (very slow with cold water, faster with steam) to Ba (vigorous with cold water).

Solubilities of the hydroxides increase down the group: Mg(OH)2 is sparingly soluble (giving the mildly basic suspension known as milk of magnesia), while Ba(OH)2 is appreciably soluble and gives a strongly basic solution. The opposite trend holds for the sulfates: BaSO4 is essentially insoluble (the basis of the BaCl2 test for sulfate ions), whereas MgSO4 is freely soluble. These trends are rationalised in terms of competition between lattice enthalpy and hydration enthalpy as the cation grows.

Thermal stabilities of the carbonates and nitrates increase down the group because the larger cation polarises the carbonate or nitrate anion less — a polarising-power argument first formalised in Fajans' rules. MgCO3 decomposes around 350 °C; BaCO3 only above 1300 °C.

RP10 asks you to verify these trends experimentally: comparing solubilities of hydroxides and sulfates, observing reactivity with water, and using flame tests (Davy's classical demonstrations) to identify Ca, Sr and Ba. See the Group 2 alkaline earth metals lesson.

Group 7: The Halogens (RP10 Anchor)

The halogens (F, Cl, Br, I) are non-metals whose chemistry is dominated by the gain of one electron to form X(-) ions or by the formation of covalent X-X and X-element bonds. Down the group, atomic radius increases, electronegativity decreases, and oxidising power decreases: F2 oxidises every other halide, Cl2 oxidises Br- and I-, Br2 oxidises only I-, and I2 oxidises none of the halides. These displacement reactions in aqueous solution are an AQA staple.

Boiling points rise down the group as the larger, more polarisable molecules experience stronger London dispersion forces (the same Drude-style induced-dipole argument used for noble gases). F2 and Cl2 are gases, Br2 is a liquid, I2 is a solid at room temperature.

The reactions of solid sodium halides with concentrated H2SO4 are a classic AQA distinction. NaF and NaCl give only the hydrogen halide (acid-base reaction, no redox). NaBr gives HBr, then some of the HBr is oxidised by H2SO4 to Br2 (with SO2 produced). NaI is the strongest reducing agent: I- reduces H2SO4 all the way to H2S (smelly), with I2 as the oxidised product. The trend in reducing power of the halides is the mirror image of the trend in oxidising power of the elements.

Aqueous chlorine disproportionates: Cl2 + H2O ⇌ HClO + HCl (used for water treatment), and Cl2 + 2NaOH → NaCl + NaClO + H2O (used to make bleach). RP10 includes the silver-nitrate / nitric-acid / ammonia sequence for distinguishing Cl-, Br- and I-. See the Group 7 halogens lesson.

Period 3 Elements and Their Oxides

The Period 3 oxides form a neat acid-base trend that AQA examines repeatedly. Na2O and MgO are basic ionic oxides that react with water (Na2O readily, MgO slowly) to give alkaline solutions. Al2O3 is amphoteric — it reacts with both acids (forming aluminium salts) and concentrated alkalis (forming aluminate ions). SiO2 is acidic but insoluble in water; it reacts only with hot concentrated alkali. P4O10 and SO3 are acidic molecular oxides that dissolve in water to give strong oxoacids: phosphoric(V) acid and sulfuric(VI) acid respectively.

The melting points of the oxides reflect their structures. Na2O and MgO are giant ionic lattices with very high melting points (MgO especially, because the Mg(2+)/O(2-) lattice is doubly charged on both ions). Al2O3 is ionic with significant covalent character (consistent with Pauling's electronegativity argument). SiO2 is a giant covalent lattice — extremely high melting. P4O10 and SO3 are molecular and melt low.

The pH of the resulting solutions traces the same acid-base pattern: Na2O gives pH ~14, MgO gives pH ~9 (limited solubility), Al2O3 and SiO2 are essentially insoluble in pure water, P4O10 gives pH ~1, SO3 gives pH ~0. A common pitfall is to forget that "acidic oxide" does not always mean "soluble" — SiO2 is acidic but inert to neutral water. See the Period 3 elements and oxides lesson.

Transition Metals: General Properties

A transition metal is, on the AQA definition, a d-block element that forms at least one stable ion with an incomplete d sub-shell. This rules out Sc (Sc(3+) is d(0)) and Zn (Zn(2+) is d(10)), even though both are d-block elements. The first-row transition metals (Ti, V, Cr, Mn, Fe, Co, Ni, Cu) are the focus.

Transition metals share four characteristic properties:

1. Variable oxidation states, because the 3d and 4s electrons are close in energy and can be removed sequentially. Mn ranges from +2 to +7; Cr from +2 to +6.
2. Formation of complex ions with ligands.
3. Coloured compounds in many oxidation states (Mn(2+) pale pink, Mn(VII) deep purple).
4. Catalytic activity, both heterogeneous (e.g. Fe in the Haber process) and homogeneous (e.g. Fe(2+)/Fe(3+) in the persulfate-iodide reaction).

The "incomplete d sub-shell" criterion is worth memorising verbatim. A common pitfall is to call Zn a transition metal because it is in the d-block. See the transition metals properties lesson.

Complex Ions and Ligands

A complex ion consists of a central metal ion bonded to surrounding species called ligands by coordinate (dative covalent) bonds — the ligand donates a lone pair to an empty orbital on the metal. The systematic chemistry of complex ions was first worked out by Werner around 1893.

Ligands are classified by the number of donor atoms (lone pairs) used to bond to the metal:

• Monodentate ligands form one bond: H2O, NH3, Cl-, CN-, OH-.
• Bidentate ligands form two bonds: ethanedioate (C2O4^2-), 1,2-diaminoethane ("en", H2NCH2CH2NH2).
• Multidentate (polydentate) ligands form many bonds: EDTA(4-) forms six bonds (hexadentate) and is widely used as a complexometric titrant.

The coordination number is the total number of coordinate bonds to the metal. Most aqueous transition-metal hexaaqua ions, e.g. Cu(H2O)6, have coordination number 6 and an octahedral shape. With larger ligands like Cl-, coordination number drops to 4 — CuCl4 is tetrahedral, while Ni(CN)4 is square planar.

Two important stereochemistry features: octahedral M-(bidentate)3 complexes are optically active (non-superimposable mirror images), and square planar M(monodentate)2(monodentate')2 complexes show cis-trans isomerism, of which cisplatin is the medicinally famous example. See the complex ions and ligands lesson.

Ligand Substitution Reactions (RP12 Anchor)

A ligand substitution reaction is the replacement of one ligand by another in a complex. AQA expects you to know three categories.

Substitution with no shape change. Cu(H2O)6 reacts with concentrated NH3 in steps: first to give [Cu(H2O)4(OH)2] (a pale-blue precipitate) and finally Cu(NH3)4(H2O)2, a deep-blue solution. Both Cu(H2O)6 and Cu(NH3)4(H2O)2 are octahedral; only four water molecules are exchanged for ammonia.

Substitution with shape change. Cu(H2O)6 reacts with concentrated HCl to give the yellow-green CuCl4 — coordination number drops from 6 (octahedral) to 4 (tetrahedral) because Cl- is larger than H2O.

Substitution with multidentate ligands. When EDTA(4-) replaces six water molecules in M(H2O)6, one EDTA molecule displaces six water molecules. The product is one complex ion plus six free water molecules — the number of free particles rises. This chelate effect is entropy-driven: ΔS is positive, so ΔG = ΔH - TΔS becomes more negative even when ΔH is nearly unchanged.

RP12 anchors these substitutions experimentally by tracking colour changes when concentrated HCl, ammonia or excess sodium hydroxide is added to aqueous copper(II), iron(II), iron(III) and cobalt(II) solutions. A common pitfall is to forget that the chelate effect is entropic — examiners want a ΔS argument, not "stronger bonds". See the ligand substitution reactions lesson.

Colour in Transition-Metal Complexes

Transition-metal complexes are coloured because d-electrons can be promoted between two sets of d orbitals split in energy by the ligand field. The splitting energy ΔE depends on the metal, its oxidation state, the ligands, and the geometry. When ΔE corresponds to a frequency in the visible region, a photon of that frequency is absorbed, an electron is promoted, and the transmitted (or reflected) light is the complementary colour. The quantitative theory of this splitting was developed by Tanabe and Sugano in the 1950s.

E = hν, so the wavelength absorbed and the colour observed depend on ΔE. Strong-field ligands (CN-, NH3) cause larger splittings and shift absorption to higher frequencies; weak-field ligands (H2O, F-) give smaller splittings. The order of ligand field strength (the spectrochemical series) is a useful AQA fact.

Three rules govern whether a complex is coloured at all:

1. The d sub-shell must be partially filled (so there are empty d orbitals to promote into). Sc(3+) (d(0)) and Zn(2+) (d(10)) are therefore colourless.
2. Octahedral and tetrahedral complexes show different splittings (and different colours) because the geometry of the ligand field differs.
3. Square planar complexes (e.g. Ni(CN)4) have their own four-orbital splitting pattern.

UV-visible spectroscopy is the analytical handle: a calibration curve of absorbance against concentration (Beer-Lambert law) lets you determine [Cu(2+)] in an unknown solution. A common pitfall is to claim that "the electron jumps to a higher shell" — the promotion is within the d sub-shell, between split d levels. See the colour in transition-metal complexes lesson.

Reactions of Aqueous Transition-Metal Ions (RP11 Anchor)

This is the largest single block of inorganic content on AQA and the one most likely to appear as a 10-12 mark Paper 1 question. The key idea is that the hexaaqua ions M(H2O)6 of transition metals are weakly acidic — they donate H+ from coordinated water — and that the acidity rises sharply with the charge on the central ion, because higher charge density polarises the O-H bond more.

The two-fold rule:

• 2+ aqua ions (e.g. Fe(H2O)6, Cu(H2O)6, Co(H2O)6) are weakly acidic. They produce neutral hydroxide precipitates with limited NaOH or NH3 (e.g. Fe(OH)2 dirty green, Cu(OH)2 pale blue).
• 3+ aqua ions (e.g. Fe(H2O)6, Al(H2O)6) are markedly more acidic — pH around 2-3. They produce hydroxide precipitates with NaOH or NH3 (Fe(OH)3 rust-brown, Al(OH)3 white).

Three reactions are examined repeatedly:

1. With OH-/NH3 (limited). Deprotonation of coordinated water to give a neutral hydroxide precipitate, e.g. Cu(H2O)6 + 2OH- → Cu(OH)2(H2O)4 + 2H2O.
2. With NaOH (excess). Amphoteric hydroxides like Al(OH)3 and Cr(OH)3 redissolve to give aluminate or chromate(III) ions; copper and iron hydroxides do not redissolve.
3. With NH3 (excess). Cu(2+) gives the deep-blue tetraamminediaqua-copper(II) complex; Cr(3+) and Co(2+) form their own ammine complexes.

A separate strand examines carbonate behaviour: 2+ aqua ions form carbonate precipitates (CuCO3, FeCO3), but 3+ aqua ions react with Na2CO3 to give the hydroxide precipitate plus CO2 — the 3+ ion is acidic enough to release CO2 from carbonate.

RP11 turns this into a wet-lab decision tree. A common pitfall is to write "Cu(2+) reacts with carbonate to give CuCO3" but then write the same for Fe(3+) — which actually gives Fe(OH)3 and CO2. See the reactions of aqueous ions lesson.

Transition-Metal Catalysis

Transition metals are exceptionally effective catalysts because their variable oxidation states and partly filled d orbitals let them form transient bonds with reactants without overcommitting — the Sabatier "just-right" binding principle.

Heterogeneous catalysis uses a solid catalyst with reactant gases or liquids. Iron is the catalyst for the Haber process (N2 + 3H2 → 2NH3); vanadium(V) oxide for the Contact process (2SO2 + O2 → 2SO3); nickel for the hydrogenation of unsaturated fats. The reaction proceeds by adsorption of reactants on active sites, surface reaction, and desorption of products. Catalysts are easily poisoned — sulfur poisons platinum catalytic converters; the lead-poisoning of the old Pt/Pd auto catalysts is the canonical example.

Homogeneous catalysis uses a catalyst in the same phase as the reactants. The classic AQA example is the iron-catalysed reaction between persulfate (S2O8(2-)) and iodide: the uncatalysed reaction is slow because two negative ions must collide. Fe(2+) provides an alternative route — Fe(2+) is oxidised to Fe(3+) by S2O8(2-), then Fe(3+) is reduced back to Fe(2+) by I-, regenerating the catalyst. Both steps involve like-charge collisions with weaker repulsion, lowering the activation energy. The autocatalysis of the MnO4- / C2O4(2-) reaction by Mn(2+) is a second AQA staple.

A common pitfall is to claim catalysts "lower activation energy by providing energy" — they provide an alternative pathway with a lower activation energy, without being consumed. See the transition-metal catalysis lesson.

Required Practicals: RP10, RP11, RP12

Three Required Practicals sit inside this block and account for a substantial fraction of Paper 3 marks:

• RP10 — Reactivity trends down Group 2 and Group 7. Reactions of Group 2 metals and oxides with water; solubility of Group 2 hydroxides and sulfates; displacement reactions between halogens and halides; the silver-nitrate test for halides; halide reactions with concentrated H2SO4.
• RP11 — Preparation and reactions of transition-metal complexes. Hexaaqua-ion reactions with limited and excess NaOH and NH3; carbonate tests; identification of unknown aqueous metal ions by colour and precipitate behaviour.
• RP12 — Investigating ligand substitution. Colour changes when concentrated HCl, ammonia or EDTA(4-) is added to aqueous Cu(2+), Co(2+) and Fe(3+) solutions; tracking shape changes from 6 to 4 coordination.

Paper 3 routinely sets six-mark "describe the procedure and predict the observation" prompts on these Required Practicals. Build a single A4 sheet per RP with the procedure, observations, equations and ionic equations, and revisit weekly.

Common Mark-Loss Patterns

• Calling Zn (or Sc) a transition metal.
• Failing to specify "incomplete d sub-shell in a stable ion".
• Forgetting that the chelate effect is entropy-driven.
• Confusing Group 2 carbonate and nitrate stability trends with the opposite direction.
• Writing "Fe(3+) + CO3(2-) → FeCO3" — it actually gives Fe(OH)3 + CO2.
• Forgetting that limited NaOH and NH3 give the same precipitate but excess gives different products.
• Forgetting the disproportionation equations for chlorine in water and alkali.
• Saying "the electron jumps to a higher energy level" without specifying within the split d orbitals.
• Reading the spectrochemical series backwards (strong-field vs weak-field).
• Forgetting that catalysts must be regenerated — every cycle must restore the original species.

How to Revise This Topic

• Build a single observation table for every Cu(2+), Fe(2+), Fe(3+), Al(3+) and Cr(3+) reaction with NaOH (limited and excess), NH3 (limited and excess), Na2CO3 and concentrated HCl. Memorise it cold.
• Draw the four shapes — octahedral, tetrahedral, square planar, linear — and label coordination numbers.
• Drill ligand substitution colour changes until you can write the formula and colour of every product in under five seconds.
• Practise displacement-equation balancing between halogens and halides, including ionic and full equations.
• Walk through the three Required Practicals with a partner, talking through procedure and expected observations.
• Use the LearningBro AI tutor on the Inorganic course to test yourself on unseen ion-identification puzzles.

Linking to Other Topics

The inorganic block touches almost every other section of AQA 7405. Lattice enthalpies for Group 2 carbonate decomposition come from energetics; the redox couples in transition-metal reactions are the E° tables from redox and electrochemistry; the acidity of 3+ aqua ions and the carbonate "release" of CO2 are direct applications of acids and buffers; and the colour-tests-and-precipitates logic carries into analytical chemistry for full ion identification. Inorganic chemistry is the synoptic glue that holds Paper 1 and Paper 3 together.

Final Word

Inorganic chemistry rewards careful pattern-matching and clean language. Master the periodicity story, the Group 2 and Group 7 trends, the Period 3 oxide acid-base scale, and the aqueous-ion decision tree, and you have built a framework that handles every unseen question AQA can set. The full LearningBro AQA A-Level Chemistry Inorganic course walks through each sub-topic with worked examples, RP10/RP11/RP12 practical anchors, and AI tutor feedback. Drill the observation tables, learn the colour changes cold, and use the synoptic links above to convert this section into a guaranteed source of Paper 1 marks.