Edexcel A-Level Chemistry: Inorganic Chemistry — Complete Revision Guide (9CH0)

Inorganic chemistry on Edexcel 9CH0 is the chemistry of the s-block, the p-block halogens, and the d-block transition metals. It looks like a long list of trends and reactions to memorise, and it is — but the trends behind the reactions are short and tightly connected to atomic structure, bonding and energetics. Once you understand why Group 2 reactivity increases down the group, why halide solubility tests work, and why transition metals form coloured complexes, the surface detail compresses into a small set of explanatory principles.

Inorganic chemistry is also one of the highest-scoring sections on Paper 1, with Group 2, Group 7 and transition-metal questions appearing reliably every series. Practical observations — colour changes in flame tests, precipitate colours in cation tests, gas evolution in carbonate decompositions — are tested directly on Paper 3 alongside the underlying explanations.

This guide is a topic-by-topic walkthrough of the inorganic content in 9CH0. It covers Group 2 reactions, solubility and thermal stability of carbonates and nitrates; Group 7 halogen reactivity, displacement reactions and disproportionation; halide ion tests with silver nitrate and ammonia; transition metal definitions, electron configurations, variable oxidation states, complex ions, ligand exchange, colour, catalysis; and qualitative analysis schemes for cation and anion identification. For each topic you will find the core ideas, common pitfalls, a worked example and a link into the LearningBro Inorganic Chemistry course.

What the Edexcel 9CH0 Specification Covers

Edexcel A-Level Chemistry (9CH0) is examined through three written papers. Paper 1 (Inorganic and Physical) is one hour 45 minutes for 90 marks and covers the inorganic content most heavily. Paper 2 (Organic and Physical) overlaps with the physical content. Paper 3 (General and Practical Principles) is two hours 30 minutes for 120 marks and tests inorganic chemistry through practical-based questions, particularly qualitative analysis.

Sub-topic	Spec area	Typical Paper 1 weight
Group 2 reactions	Topic 4	4-6 marks
Solubility of Group 2 compounds	Topic 4	3-5 marks
Thermal stability of carbonates and nitrates	Topic 4	3-5 marks
Group 7 reactivity and displacement	Topic 4	4-6 marks
Halide ion tests (AgNO3 / NH3)	Topic 4	3-5 marks
Concentrated H2SO4 with halides	Topic 4	4-6 marks
Transition metal definitions and configurations	Topic 15	3-5 marks
Variable oxidation states and complexes	Topic 15	6-10 marks
Colour and ligand exchange	Topic 15	4-6 marks
Catalysis (homogeneous and heterogeneous)	Topic 15	3-5 marks
Qualitative analysis schemes	Paper 3	6-10 marks

These weights are estimates, modelled on the Edexcel 9CH0 paper format. What is reliable is that a transition-metal extended-response question — usually colour or ligand exchange — appears on essentially every Paper 1, and a qualitative-analysis sequence appears on essentially every Paper 3.

Group 2: Reactions, Solubility and Thermal Stability

Group 2 metals (Be, Mg, Ca, Sr, Ba) become more reactive down the group because the outer 2 electrons are progressively further from the nucleus and shielded by more inner shells, so they are easier to remove. Reactions with water show the trend clearly: Mg reacts very slowly with cold water and readily with steam to give MgO + H2; Ca reacts noticeably with cold water to give Ca(OH)2 + H2; Sr more vigorously; Ba very vigorously. Reactions with oxygen all give MO with progressively more vigorous combustion. Reactions with dilute acids give the chloride or sulfate plus hydrogen, again accelerating down the group.

Solubility of Group 2 compounds shows two opposite trends. The hydroxides become more soluble down the group (Mg(OH)2 is almost insoluble; Ba(OH)2 is freely soluble). The sulfates become less soluble down the group (MgSO4 is freely soluble; BaSO4 is essentially insoluble — hence its use as the standard sulfate-ion test reagent). The qualitative explanation involves competition between lattice enthalpies (which become less exothermic down the group as ions get larger) and hydration enthalpies (which also become less exothermic, but at a different rate); the relative changes determine the direction of the solubility trend.

Thermal stability of Group 2 carbonates and nitrates increases down the group. MgCO3 decomposes at around 350 °C; BaCO3 requires over 1300 °C. The reason is the polarising power of the cation: small, highly charged cations (like Mg^2+) polarise the carbonate ion strongly, weakening one of the C-O bonds and making decomposition easier. As cation size increases down the group, polarising power decreases and decomposition needs more heat.

Compound	Decomposition temperature (°C, approx.)	Trend
MgCO3	350	Easiest to decompose
CaCO3	825
SrCO3	1280
BaCO3	1360	Hardest to decompose

Worked example. Predict whether Mg(NO3)2 or Ba(NO3)2 has the higher thermal stability. Ba^2+ is larger than Mg^2+ and has lower polarising power, so it polarises the nitrate ion less strongly. Ba(NO3)2 is therefore more thermally stable. In fact, all Group 2 nitrates decompose to MO + NO2 + O2 on heating.

A common pitfall is to swap the trends in hydroxide and sulfate solubility. A useful mnemonic is "hydroxides up, sulfates down" — solubility of hydroxides increases, sulfates decrease, both moving down the group. See the Group 2 lesson for full reaction equations and observations.

Group 7: Reactivity, Displacement and Disproportionation

Group 7 halogens (F, Cl, Br, I) become less reactive down the group as oxidising agents because the outer p electron-accepting orbital is further from the nucleus and more shielded, making it harder to attract an electron to form a halide. Reactivity falls F > Cl > Br > I.

Displacement reactions are the standard demonstration. A more reactive halogen (higher in the group) displaces a less reactive one from solution: Cl2 + 2KBr → 2KCl + Br2, observable as the colourless solution becoming orange. Br2 + 2KI → 2KBr + I2 turns the solution brown. Iodine cannot displace bromine, so I2 + KBr gives no reaction. The colours in cyclohexane (or another non-polar solvent) sharpen the test: chlorine pale green, bromine orange, iodine purple.

Disproportionation is when one species is simultaneously oxidised and reduced. The classic Group 7 example is chlorine in water and in cold and hot alkali:

• Cl2 + H2O ⇌ HCl + HClO (cold; produces bleach)
• Cl2 + 2NaOH → NaCl + NaClO + H2O (cold dilute; bleach manufacture)
• 3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O (hot concentrated)

In each, Cl goes from oxidation state 0 to both -1 (in chloride) and either +1 (in hypochlorite) or +5 (in chlorate). Working out half-equations and assigning oxidation states is a frequent question.

Worked example. Identify which atom is oxidised and which is reduced in 3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O. In Cl2, Cl is 0. In NaCl, Cl is -1 (reduced). In NaClO3, Cl is +5 (oxidised). The same Cl element is both oxidised and reduced — this is the definition of disproportionation.

A common pitfall is to call any redox reaction "disproportionation" — it must be the same element undergoing both processes. See the Group 7 lesson for full equation drills.

Halide Ion Tests

The standard test for halide ions in solution uses silver nitrate and ammonia in sequence. Add dilute nitric acid first (to remove any carbonate or hydroxide that would also precipitate with silver), then add silver nitrate. A precipitate forms; its colour and solubility in ammonia distinguish the halide.

Halide	AgX precipitate colour	Dilute NH3	Concentrated NH3
Cl-	White	Dissolves	Dissolves
Br-	Cream	Insoluble	Dissolves
I-	Yellow	Insoluble	Insoluble

The pattern reflects the strength of the Ag-X bond: AgCl has the weakest bond and is most easily broken by ammonia ligand substitution to form [Ag(NH3)2]+; AgI has the strongest bond and is not displaced.

A second halide test is the reaction with concentrated sulfuric acid on the solid sodium halide. NaCl gives only HCl gas (steamy fumes, white smoke with NH3). NaBr gives HBr plus some Br2 (brown fumes) because HBr is partly oxidised by H2SO4. NaI gives a complex mixture: HI, I2 (purple fumes/black solid), H2S (rotten-egg smell), S (yellow solid) — H2SO4 is strongly oxidising and is reduced progressively further by the more strongly reducing iodide. The trend confirms the order of reducing power: I- > Br- > Cl-.

A common pitfall is to forget the dilute nitric acid pre-treatment, or to mix up which halide dissolves in which strength of ammonia. See the halide tests lesson for full observation tables.

Transition Metals: Definitions and Configurations

A transition metal is a d-block element that forms at least one stable ion with a partially filled d sub-shell. By this definition, scandium (only Sc^3+, d^0) and zinc (Zn^2+, d^10) are not transition metals although they sit in the d-block. The other first-row d-block elements (Ti to Cu) are.

Transition metals show four characteristic properties: variable oxidation states, coloured ions, complex ion formation and catalytic activity. All four follow from the partially filled d sub-shell. Their electron configurations follow the pattern [Ar] 3d^n 4s^2 with the chromium and copper anomalies (Cr is [Ar] 3d^5 4s^1, Cu is [Ar] 3d^10 4s^1). When ionising, the 4s electrons leave first.

Worked example. Write the electron configuration of Mn, Mn^2+ and Mn^7+ (the formal oxidation state in MnO4-). Mn is [Ar] 3d^5 4s^2. Mn^2+ is [Ar] 3d^5 — half-filled and stable. Mn^7+ would be [Ar] (no d electrons left, but the +7 species exists only in covalent compounds like MnO4-, where the bonding is largely covalent and the formal oxidation state is +7).

A common pitfall is to remove 3d electrons before 4s — the order is reversed on ionisation. Another is to call zinc a transition metal. See the transition metal lesson.

Complex Ions and Ligand Exchange

A complex ion is a transition metal ion bonded by dative covalent bonds to one or more ligands (electron-pair donors). Common ligands are H2O, NH3 (both monodentate: one donor atom), Cl-, OH-, CN-, and bidentate ligands such as ethanedioate (C2O4^2-) and ethane-1,2-diamine (en). EDTA^4- is a hexadentate ligand. The total number of dative bonds is the coordination number.

Six-coordinate complexes are octahedral (e.g. [Cu(H2O)6]^2+). Four-coordinate complexes are usually tetrahedral (e.g. [CuCl4]^2-, smaller ligands and steric crowding) or square planar (less common, [PtCl4]^2-). Two-coordinate complexes are linear (e.g. [Ag(NH3)2]+).

Ligand exchange reactions swap one ligand for another. The colour change accompanying the substitution is a frequent observation question. Adding excess ammonia to copper(II) sulfate, for example, replaces four of the six water ligands to give [Cu(NH3)4(H2O)2]^2+, with the colour changing from pale blue to deep royal blue. Adding concentrated HCl to copper(II) sulfate replaces all six water ligands with four chloride ligands, giving the tetrahedral [CuCl4]^2-, which is yellow-green.

Complex	Colour	Geometry
[Cu(H2O)6]^2+	Pale blue	Octahedral
[Cu(NH3)4(H2O)2]^2+	Deep blue	Octahedral
[CuCl4]^2-	Yellow-green	Tetrahedral
[Fe(H2O)6]^2+	Pale green	Octahedral
[Fe(H2O)6]^3+	Pale violet (often yellow due to hydrolysis)	Octahedral
[Ag(NH3)2]+	Colourless	Linear

Worked example. Write equations for the stepwise addition of NH3(aq) to [Cu(H2O)6]^2+. (1) Adding small amounts of NH3 first deprotonates two water ligands to give the pale blue precipitate Cu(OH)2(H2O)4. (2) Excess NH3 dissolves the precipitate by ligand substitution: Cu(OH)2(H2O)4 + 4NH3 → [Cu(NH3)4(H2O)2]^2+ + 2OH- + 2H2O.

A common pitfall is to confuse the precipitate (Cu(OH)2) and complex stages, or to swap the colours. See the complex ions lesson for full colour and equation tables.

Colour and the d-Orbital Splitting

Transition metal complex ions are coloured because the d orbitals are split into two groups by the ligand field, and electrons can be promoted from the lower group to the higher group by absorbing a photon of visible light. The colour observed is the complementary colour of the wavelength absorbed.

The size of the d-orbital splitting (and therefore the wavelength absorbed) depends on the metal, the oxidation state, the ligand and the geometry. The spectrochemical series orders ligands by their ability to split the d orbitals: I- < Br- < Cl- < OH- < H2O < NH3 < CN-. So changing ligand changes the absorbed wavelength and hence the colour.

Sc^3+ (d^0) and Zn^2+ (d^10) are colourless because there are no d-d transitions available. Most other transition metal ions are coloured.

A common pitfall is to attribute colour to "the d orbitals being half full" — the explanation is d-orbital splitting and absorption of visible light. See the colour and splitting lesson.

Catalysis (Homogeneous and Heterogeneous)

Transition metals catalyse a wide range of reactions because they can adopt variable oxidation states and provide alternative reaction pathways with lower activation energies.

Heterogeneous catalysts are in a different phase from the reactants. Examples on the specification include iron in the Haber process (N2 + 3H2 → 2NH3), V2O5 in the Contact process (2SO2 + O2 → 2SO3) and platinum/rhodium in catalytic converters (CO + NO + hydrocarbons → CO2 + N2 + H2O). The mechanism is adsorption of reactants onto active sites on the catalyst surface, weakening internal bonds, allowing reaction, then desorption of products.

Homogeneous catalysts are in the same phase as the reactants, typically aqueous. The classic example is the catalysis of S2O8^2- + 2I- → 2SO4^2- + I2 by Fe^2+ ions: the Fe^2+ is oxidised to Fe^3+ by S2O8^2-, then reduced back to Fe^2+ by I-, regenerating the catalyst. The two-step pathway has a lower activation energy than direct reaction between two negative ions.

A common pitfall is to write a catalyst as being "consumed" — it is regenerated by the cycle. See the catalysis lesson for the full mechanism diagrams.

Qualitative Analysis: Cation and Anion Tests

Paper 3 includes a qualitative-analysis sequence: identify ions in an unknown solution from a series of test results. The standard tests are summarised below.

Ion	Reagent	Observation
Cu^2+	NaOH(aq)	Pale blue precipitate; insoluble in excess
Fe^2+	NaOH(aq)	Green precipitate; turns brown on standing
Fe^3+	NaOH(aq)	Brown precipitate
Al^3+	NaOH(aq)	White precipitate; dissolves in excess (amphoteric)
NH4+	NaOH then warm	Ammonia gas evolved (turns damp red litmus blue)
CO3^2-	Dilute HCl	CO2 evolved (turns limewater milky)
SO4^2-	BaCl2 in dilute HCl	White precipitate of BaSO4
Cl-	AgNO3 in dilute HNO3	White precipitate; dissolves in dilute NH3
Br-	AgNO3 in dilute HNO3	Cream precipitate; dissolves in conc NH3 only
I-	AgNO3 in dilute HNO3	Yellow precipitate; insoluble in conc NH3

A common pitfall is to forget the order of testing — anions tested with silver nitrate must first have any carbonate removed by dilute nitric acid, or false positive precipitates form. Always acidify before precipitation tests. See the qualitative analysis lesson.

Common Mark-Loss Patterns

• Confusing the directions of Group 2 hydroxide and sulfate solubility trends.
• Calling any redox reaction "disproportionation".
• Removing 3d electrons before 4s when ionising a transition metal.
• Calling Sc and Zn "transition metals" — they are not.
• Stating colours without writing the geometry, or geometry without colours.
• Forgetting the dilute HNO3 pre-treatment in halide tests.
• Calling a catalyst "consumed".
• Writing "ligand exchange" without identifying which ligand replaces which.
• Failing to balance complex-ion charge in equations.
• Forgetting state symbols, especially (g) for evolved gases like CO2 and NH3.

How to Revise This Topic

• Build a single sheet of equations and observations for Group 2, Group 7 and the halide tests. Recite it from memory three times a week.
• Make a colour deck of every transition-metal complex on the specification, with formula on one side and colour and geometry on the other. Drill until automatic.
• Practise qualitative-analysis sequences under timed conditions. Examiners expect you to deduce the ion in two or three sentences.
• Drill ligand-exchange equations for [Cu(H2O)6]^2+ with NH3, Cl- and OH-. These are the most-examined complex-ion reactions.
• Use the LearningBro practice quizzes to test recall under exam conditions.

Linking to Other Topics

Inorganic chemistry is heavily synoptic. Group 2 and Group 7 reactivity trends are direct applications of the atomic structure trends. Lattice and hydration enthalpy reasoning underlies Group 2 solubility (see energetics). Halide test reactions are applications of bonding and complex-ion formation. Transition-metal redox reactions reuse the half-equation framework from amounts and redox, and many qualitative tests double as questions in acids and buffers.

Final Word

Inorganic chemistry rewards memorisation paired with explanation. Learn the colours, equations and observations cold, but always be ready to explain why each trend goes the way it does. The full LearningBro Inorganic Chemistry course walks through every topic with diagrams, observation tables and AI tutor feedback. Get this section right and Paper 1 becomes much shorter and Paper 3's qualitative analysis becomes a routine exercise.