AQA A-Level Chemistry: Inorganic Chemistry and Analytical Chemistry Guide

Inorganic chemistry and analytical chemistry are two pillars of the AQA A-Level Chemistry specification that often receive less revision attention than organic chemistry, yet they carry substantial weight in the exams. Inorganic chemistry demands a strong command of periodic trends, group properties, and the behaviour of transition metals. Analytical chemistry requires you to interpret spectra, identify unknown compounds, and apply a range of instrumental techniques with precision.

This guide covers the AQA specification content for both areas -- from periodicity and the reactions of Period 3 elements through to NMR spectroscopy and chromatography -- giving you a thorough foundation for exam success.

Inorganic Chemistry

Periodicity

Periodicity is the study of recurring trends in element properties as you move across a period or down a group. You need to understand and explain trends in atomic radius, first ionisation energy, electronegativity, and melting point across Periods 2 and 3.

Atomic radius decreases across a period because nuclear charge increases while electrons are added to the same principal energy level. Down a group, atomic radius increases because additional principal energy levels and increased shielding outweigh the higher nuclear charge.

First ionisation energy generally increases across a period. However, there are two key discontinuities. Between Groups 2 and 3, aluminium has a lower first ionisation energy than magnesium because its outer electron occupies a higher-energy 3p orbital. Between Groups 5 and 6, oxygen has a lower first ionisation energy than nitrogen because the fourth 2p electron must pair, creating inter-electron repulsion that makes removal easier.

Electronegativity increases across a period and decreases down a group, following the same reasoning related to nuclear charge, atomic radius, and shielding.

Melting points across Period 3 show a distinctive pattern. Sodium, magnesium, and aluminium have increasingly strong metallic bonding as the charge on the metal ion increases and the ionic radius decreases. Silicon has a very high melting point due to its giant covalent (macromolecular) structure. Phosphorus, sulphur, chlorine, and argon form simple molecular structures with low melting points governed by weak intermolecular forces -- sulphur is higher than phosphorus because S8 molecules experience stronger London dispersion forces than P4.

Properties of Period 3 Elements and Their Oxides

The oxides of Period 3 elements show a systematic change in acid-base character across the period. Sodium oxide and magnesium oxide are basic -- they dissolve in water to give alkaline solutions. Aluminium oxide is amphoteric, reacting with both acids and bases. Silicon dioxide is weakly acidic with a giant covalent structure. Phosphorus(V) oxide and sulphur trioxide are strongly acidic, dissolving in water to form phosphoric acid and sulphuric acid respectively.

The trend is driven by the change from ionic bonding on the left of the period to covalent bonding on the right. You should be able to write equations for the reactions of each oxide with water and explain the acid-base character in terms of bonding and structure.

Group 2 -- The Alkaline Earth Metals

Reactivity with water increases down the group. Magnesium reacts slowly with cold water but vigorously with steam, while calcium, strontium, and barium react increasingly vigorously with cold water. This is explained by decreasing ionisation energy down the group -- increased shielding and atomic radius mean outer electrons are lost more easily.

Solubility of hydroxides increases down Group 2, from sparingly soluble magnesium hydroxide to freely soluble barium hydroxide. Solubility of sulphates follows the opposite trend -- magnesium sulphate is soluble while barium sulphate is insoluble, which forms the basis of the barium chloride test for sulphate ions.

Thermal stability of carbonates and nitrates increases down the group. Smaller cations have greater polarising power, distorting the anion electron cloud more effectively and making decomposition easier. This is why magnesium carbonate decomposes at a lower temperature than barium carbonate.

Group 7 -- The Halogens

Electronegativity and oxidising power decrease down Group 7. Chlorine can displace bromine and iodine from solution, but bromine can only displace iodine. You should be confident writing ionic equations for these displacement reactions.

Halide ion tests with silver nitrate produce precipitates of different colours -- white for chloride, cream for bromide, and yellow for iodide. Silver chloride dissolves in dilute ammonia, silver bromide in concentrated ammonia, and silver iodide is insoluble in ammonia.

Reactions with concentrated sulphuric acid reveal increasing reducing power down the group. Chloride ions only produce HCl (steamy white fumes). Bromide ions reduce the acid to SO2, also producing bromine (orange-brown fumes). Iodide ions reduce it further to sulphur and hydrogen sulphide, producing iodine vapour, a yellow solid, and the smell of rotten eggs.

Disproportionation of chlorine with cold dilute NaOH produces NaCl and NaClO, while hot concentrated NaOH produces NaCl and NaClO3. You should assign oxidation states to demonstrate that chlorine is simultaneously oxidised and reduced.

Transition Metals

Transition metals are d-block elements that form at least one stable ion with a partially filled d sub-shell -- this excludes scandium and zinc.

Variable oxidation states arise because the 3d and 4s electrons are close in energy. Iron exists as Fe2+ and Fe3+, copper as Cu+ and Cu2+, and manganese across a wide range from +2 to +7.

Complex ions form when transition metal ions accept lone pairs from ligands such as water, ammonia, chloride, and cyanide. Coordination numbers are typically 6 (octahedral) or 4 (tetrahedral or square planar). Know the formulae of key complexes including [Cu(H2O)6]2+, [CuCl4]2-, and [Cu(NH3)4(H2O)2]2+.

Colour arises from d-d transitions -- ligands cause d-orbital splitting, and electrons absorb visible light when jumping between energy levels. Changes in oxidation state, ligand, or coordination number alter the splitting and therefore the colour.

Catalytic activity results from the ability to change oxidation state, providing alternative reaction pathways with lower activation energies. Iron catalyses the Haber process, vanadium(V) oxide catalyses the Contact process, and manganese(IV) oxide catalyses the decomposition of hydrogen peroxide. Transition metals also act as heterogeneous catalysts by providing a surface on which reactant molecules can adsorb and react.

Reactions of Ions in Aqueous Solution

You must know the observations when NaOH and ammonia are added to solutions of common transition metal ions. Sodium hydroxide produces a green precipitate with Fe2+ (darkens on standing), a brown precipitate with Fe3+, and a blue precipitate with Cu2+.

In excess ammonia, the copper(II) hydroxide precipitate dissolves to form the deep blue [Cu(NH3)4(H2O)2]2+ complex -- a ligand exchange reaction. Adding concentrated HCl to [Cu(H2O)6]2+ causes a colour change from blue to yellow-green as [CuCl4]2- forms.

Analytical Chemistry

Tests for Ions

Cation tests include the NaOH precipitation tests for transition metal ions and flame tests for Group 1 and 2 metals -- lithium (crimson), sodium (yellow), potassium (lilac), calcium (orange-red), and barium (green). Ammonium ions are identified by warming with NaOH and testing for ammonia gas with damp red litmus paper.

Anion tests include the silver nitrate test for halides, the barium chloride test for sulphates (white precipitate), and the acid test for carbonates (effervescence with CO2 that turns limewater milky).

Mass Spectrometry

Mass spectrometry determines relative atomic masses and molecular formulae by ionising samples and separating ions by their mass-to-charge ratio (m/z). For elements, isotope peaks allow you to calculate relative atomic mass using a weighted mean. For compounds, the molecular ion peak (M+) gives the relative molecular mass, while fragmentation patterns reveal structural features. Learn common fragment losses -- 15 (CH3), 17 (OH), 29 (CHO), and 45 (OC2H5).

Infrared Spectroscopy

IR spectroscopy identifies functional groups by measuring absorption of infrared radiation at characteristic wavenumbers. Key absorptions include the broad O-H stretch in alcohols (around 3230--3550 cm-1), the broader O-H in carboxylic acids (around 2500--3300 cm-1), the sharp C=O stretch (around 1680--1750 cm-1), and N-H in amines (around 3300--3500 cm-1). You receive a data sheet in the exam, but familiarity with these approximate positions and whether absorptions appear broad or sharp will help you work efficiently.

A common exam task is distinguishing between structures using IR -- for example, identifying a carboxylic acid by its broad O-H alongside a C=O absorption, which an aldehyde would lack. Another frequent question involves tracking a reaction by comparing the IR spectra of reactant and product -- for instance, confirming that an alcohol has been oxidised to a carboxylic acid by observing the disappearance of the alcohol O-H and the appearance of the broader carboxylic acid O-H and C=O absorptions.

NMR Spectroscopy

Carbon-13 NMR shows the number of distinct carbon environments in a molecule. The chemical shift indicates the type of environment, and equivalent carbons produce a single peak.

Proton (1H) NMR provides richer information: the number of hydrogen environments (number of peaks), relative numbers of hydrogens (integration), environment type (chemical shift), and neighbouring hydrogens (splitting pattern). The n+1 rule determines splitting -- a signal split into a triplet, for example, has two non-equivalent hydrogens on adjacent carbons.

TMS (tetramethylsilane) is the reference standard at 0 ppm, chosen because it gives a single sharp peak away from most organic signals and is inert and volatile. D2O shakes identify O-H and N-H protons -- any peak that disappears after adding deuterium oxide must be due to a labile proton.

Chromatography

Chromatography separates mixtures based on differential distribution between a stationary phase and a mobile phase. Thin-layer chromatography (TLC) uses a solid stationary phase on a plate, with components identified by Rf values. It is useful for checking purity and monitoring reactions.

Gas chromatography (GC) separates volatile components by retention time as they pass through a column. Peak areas are proportional to the amount of each component. GC-MS couples separation with mass spectrometric identification for powerful analysis of complex mixtures.

Column chromatography uses the same principles as TLC but on a preparative scale, with a column packed with the stationary phase to separate and collect individual components of a mixture.

The general principle underlying all chromatography is the same: components that interact more strongly with the stationary phase move more slowly and are retained longer, while those with greater affinity for the mobile phase move faster. Understanding this principle allows you to reason about any chromatographic system, even one you have not encountered before.

Exam Strategy

For inorganic chemistry, practise active recall of colours, observations, and trends regularly. Always explain the underlying cause of a trend -- stating that reactivity increases is not enough without explaining why. Ensure you can write balanced equations and ionic equations for every key reaction.

For analytical chemistry, be systematic with spectral interpretation. In NMR, note the number of peaks, integration, chemical shifts, and splitting patterns before proposing a structure. In IR, focus on diagnostic absorptions above 1500 cm-1 rather than the fingerprint region. In mass spectrometry, find the molecular ion peak first, then identify key fragment losses.

Prepare with LearningBro

LearningBro offers structured courses that cover these topics through active recall questions and spaced repetition, helping you build the knowledge and confidence needed for exam success:

• AQA A-Level Chemistry: Inorganic Chemistry in Depth -- comprehensive coverage of periodicity, Group 2, Group 7, transition metals, and Period 3 properties.
• AQA A-Level Chemistry: Inorganic Chemistry -- Acids and Reactions -- focused practice on acid-base reactions, ionic equations, and the reactions of ions in aqueous solution.
• AQA A-Level Chemistry: Analytical Chemistry in Depth -- detailed practice covering mass spectrometry, IR spectroscopy, NMR spectroscopy, chromatography, and tests for ions.

Related Reading

• A-Level Chemistry Revision Guide -- a broader guide covering organic, inorganic, and physical chemistry alongside exam technique.
• Spaced Repetition: The Science Behind Effective Revision -- understand why spaced repetition is the most efficient way to memorise the factual content that inorganic and analytical chemistry demand.
• A-Level Revision Strategy: From Mocks to Finals -- a structured plan for the final stretch of revision before your exams.

Final Thoughts

Inorganic chemistry and analytical chemistry reward thorough preparation. The factual content in inorganic chemistry -- colours, observations, trends, and equations -- must be committed to memory through repeated active recall. Analytical chemistry requires you to practise interpreting spectra and data until the process feels intuitive rather than daunting.

Do not leave these topics until the last minute. Build them into your revision schedule early, revisit them regularly, and test yourself under exam conditions with past papers. The students who perform best in these areas are not necessarily the most naturally talented -- they are the ones who have practised the most consistently.

Start with the fundamentals, build up to more demanding problems, and trust that every practice question brings you closer to the grade you are working towards.