OCR A-Level Chemistry: Periodicity, Group 2 and Halogens — Complete Revision Guide (H432)
OCR A-Level Chemistry: Periodicity, Group 2 and Halogens
Periodicity is the integrating idea of inorganic chemistry. Once you can write an electron configuration and assign a bond type from Acids, Redox, Electrons and Bonding, the periodic table stops being a memorisation chart and becomes a prediction engine. OCR A-Level Chemistry A (H432) leans on periodicity throughout: the s, p, d, f block divisions justify why transition metals behave differently from main-group elements, the ionisation-energy graph for Period 3 is a Paper 1 fixture, the descriptive chemistry of Group 2 and the halogens recurs on every paper, and the qualitative anion tests anchor a reliable practical-skills item.
H432 examiners weight this module heavily because the periodic trends provide a quantitative anchor for descriptive chemistry that would otherwise feel arbitrary. Why is sodium less reactive than potassium? The first ionisation energy is lower for potassium because the outer electron is further from the nucleus and more shielded. Why does the solubility of Group 2 sulfates decrease down the group while the hydroxides increase? Because lattice and hydration enthalpies scale differently with cation size — a calculation made fully quantitative in energetics and electrode potentials. Candidates who internalise the trends as the consequence of four named factors (nuclear charge, distance, shielding, spin-pair repulsion) find every Period 3 ionisation-energy or Group 2 reactivity question collapses into a chain of those four factors. Candidates who treat the trends as facts to memorise struggle on the explanation items where the underlying reasoning is the mark-bearing step.
Course 3 of the H432 Chemistry learning path on LearningBro, Periodicity, Group 2 and Halogens, develops the periodic-table reasoning the rest of the path will reuse. It opens with the s/p/d block organisation, develops the atomic-radius and ionisation-energy trends and reads them off the standard graphs, then walks through the structure and bonding of the Period 3 elements, the reactivity trends of Group 2 metals and their oxides/hydroxides/carbonates, the oxidising trend of the halogens and the reducing trend of their halides, and closes with the standard qualitative anion tests. It sits adjacent to Acids, Redox, Electrons and Bonding and feeds into Enthalpy, Rates and Equilibrium, Transition Elements and Aromatic and downstream into every redox context on the OCR A-Level Chemistry learning path.
Guide Overview
The Periodicity, Group 2 and Halogens course is built as a sequence of lessons that move from periodic organisation through trend analysis into the descriptive chemistry of Group 2 and the halogens.
- The Periodic Table and Blocks
- Atomic Radius and Ionisation Energy Trends
- Successive Ionisation Energies
- Period 3 Structure and Bonding
- Group 2 Reactivity Trends
- Group 2 Compounds and Applications
- Halogen Oxidising Trends and Disproportionation
- Halide Reducing Trends with Sulfuric Acid
- Qualitative Anion Analysis
OCR H432 Specification Coverage
This course addresses OCR H432 Module 3.1.1 (periodicity), Module 3.1.2 (Group 2) and Module 3.1.3 (the halogens). The specification organises the topic into periodic trends, the descriptive chemistry of two representative groups, and the qualitative tests for anions that recur across practical chemistry (refer to the official OCR specification document for exact wording).
| Sub-topic | Spec area | Primary lesson(s) |
|---|---|---|
| The periodic table; s, p, d, f blocks | OCR H432 Module 3.1.1 | The Periodic Table and Blocks |
| Atomic radius, first ionisation energy and the underlying factors | OCR H432 Module 3.1.1 | Atomic Radius and Ionisation Energy Trends; Successive Ionisation Energies |
| Period 3 element structure, bonding and physical properties | OCR H432 Module 3.1.1 | Period 3 Structure and Bonding |
| Group 2 reactions with water, oxygen and acids; oxides and hydroxides | OCR H432 Module 3.1.2 | Group 2 Reactivity Trends; Group 2 Compounds and Applications |
| Halogen oxidising power and disproportionation reactions | OCR H432 Module 3.1.3 | Halogen Oxidising Trends and Disproportionation |
| Halide reducing power with concentrated sulfuric acid | OCR H432 Module 3.1.3 | Halide Reducing Trends with Sulfuric Acid |
| Tests for halide, carbonate and sulfate anions | OCR H432 Module 3.1.3 | Qualitative Anion Analysis |
Module 3.1 is examined on Paper 1 (Periodic Table, Elements and Physical Chemistry) and synoptically on Paper 3. The Period 3 ionisation-energy graph and the qualitative anion test sequence are particularly reliable mark-bearing items.
Topic-by-Topic Walkthrough
The Periodic Table and Blocks
The periodic table and blocks lesson develops the four block divisions according to the highest-energy electron's sublevel: s-block (Groups 1 and 2, with one or two outer s electrons), p-block (Groups 13-18, filling 2p through 6p), d-block (transition metals, filling 3d through 5d, characterised in detail in transition elements) and f-block (lanthanides and actinides, off the main grid). This block structure recovers the descriptive chemistry from the electron configuration, and it justifies why Group 1 and Group 2 are reducing metals (low ionisation energy) and why Group 17 are oxidising non-metals (high electronegativity, easy to add an electron).
Atomic Radius and Ionisation Energy Trends
The atomic radius and ionisation energy lesson develops the four controlling factors of first ionisation energy: nuclear charge, distance from the nucleus, shielding by inner electrons and the spin-pairing penalty within a sublevel. Atomic radius decreases across a period (constant shells, increasing nuclear charge pulls electrons closer) and increases down a group (additional shells outweigh the increasing nuclear charge because of inner-electron shielding). First ionisation energy mirrors this: it increases across a period and decreases down a group, with two characteristic discontinuities — a drop between Be and B (2p sublevel begins, 2p slightly higher energy than 2s and slightly better shielded) and a drop between N and O (the first electron pair in a 2p orbital introduces spin-pair repulsion). These two discontinuities are reliable Paper 1 explanation items.
The successive ionisation energies lesson develops the staircase pattern when an atom is stripped electron by electron. The jumps in successive IE reveal the shell structure: a sharp jump signals that the next electron is from an inner, more strongly held shell. The classic worked example is sodium — IE1 is low, then IE2 through IE9 are progressively larger but smooth, then a huge jump to IE10 and IE11 because those electrons are 1s. The number of small steps before the first big jump identifies the group number.
Period 3 Structure and Bonding
The Period 3 lesson walks through the structure, bonding and melting point of Na, Mg, Al, Si, P, S, Cl and Ar. The first three are metallic (melting point rises with cation charge and delocalised electron density). Silicon is giant covalent and has the highest melting point of the period. Phosphorus (P₄), sulfur (S₈) and chlorine (Cl₂) are simple molecular and have low melting points, with sulfur higher than phosphorus higher than chlorine because of the molecular size and consequent dispersion forces. Argon is monatomic. The Paper 1 question is to plot or interpret a melting-point bar chart and justify the trend in terms of structure and bonding, which connects directly back to bonding.
Group 2 Reactivity and Compounds
The Group 2 reactivity lesson develops the increasing reactivity down the group: Be is unreactive with water, Mg reacts slowly with cold water but vigorously with steam to give MgO, Ca reacts with cold water to give Ca(OH)₂ and H₂, Sr and Ba react more vigorously still. The trend traces back to ionisation energy — Group 2 reactivity correlates with how easily the M²⁺ ion forms. Reactions with oxygen give MO (white solid), with chlorine give MCl₂, with dilute acids give the salt + H₂. The Group 2 compounds lesson covers the descriptive chemistry of the oxides and hydroxides (basic, Mg(OH)₂ used as an antacid, Ca(OH)₂ as a soil neutraliser) and the trend in solubility of the hydroxides (increases down the group) and the sulfates (decreases down the group, with BaSO₄ insoluble — the basis of the sulfate test in qualitative analysis).
Halogen Oxidising and Halide Reducing Trends
The halogen oxidising trends lesson develops the decreasing oxidising power down the group: F₂ > Cl₂ > Br₂ > I₂, with the displacement-reaction logic that a more reactive halogen oxidises the halide ion of a less reactive halogen. Worked example: Cl₂ + 2KBr → 2KCl + Br₂ (chlorine displaces bromide) gives an orange aqueous layer, distinguishable from the brown of I₂. Disproportionation reactions cover Cl₂ + 2NaOH → NaCl + NaClO + H₂O (cold dilute base, the chemistry of household bleach) and Cl₂ + H₂O → HCl + HClO (water treatment, the basis of swimming-pool sanitation). The halide reducing trends lesson develops the converse trend: I⁻ > Br⁻ > Cl⁻ > F⁻ as reducing agents, traced via the increasingly extensive reduction of concentrated H₂SO₄ — chloride only gives HCl + steamy fumes, bromide reduces H₂SO₄ to SO₂, iodide reduces it all the way to H₂S with intermediate steps producing SO₂ and S.
Qualitative Anion Analysis
The qualitative anion analysis lesson covers the standard three-test sequence: test for carbonate first (effervescence with dilute HCl, gas turns limewater milky), then for sulfate (add acidified BaCl₂, white precipitate of BaSO₄ confirms), then for halides (add acidified AgNO₃ — white precipitate of AgCl soluble in dilute NH₃, cream AgBr soluble in concentrated NH₃, yellow AgI insoluble in concentrated NH₃). The order matters — carbonate first to avoid false positives with the BaCl₂ test (carbonate also gives a white precipitate with Ba²⁺ but it would dissolve in the acidification step). The test sequence anchors PAG 4 (Qualitative analysis).
A Typical H432 Paper 1 Question
A standard Paper 1 prompt gives candidates the first-ionisation-energy plot for Period 3 (Na to Ar) and asks them to (a) describe the overall trend, (b) explain the two characteristic discontinuities, and (c) account for the position of one specified element. The route is fixed: state the overall increase across the period (rising effective nuclear charge with no extra shell, electrons held more tightly); identify the dip between Mg and Al as the 3s-to-3p sublevel transition (3p slightly higher energy, slightly better shielded); identify the dip between P and S as the spin-pair repulsion that arises when the first 3p electron pair forms; then for the specified element, give the configuration and identify which factor dominates. The discriminator at the top band is the explicit use of the word "shielding" alongside "nuclear charge" rather than "attraction to the nucleus", and the explicit recognition that the dip at S does not reverse the overall trend — sulfur's IE1 is still higher than phosphorus's neighbour-to-the-left silicon, just lower than phosphorus's.
Worked Examples You Should Be Able to Reproduce
The fastest way to make this module automatic is to rehearse the handful of calculations and explanation-chains that recur across every H432 series. Work through each of the following with a pen before you read the model reasoning — the goal is to be able to generate the answer, not merely recognise it.
Reading successive ionisation energies to identify a group
A common Paper 1 item gives a table of successive ionisation energies (in kJ mol⁻¹) for an unknown element and asks you to identify the group it belongs to. Suppose the successive values are approximately 738, 1451, 7733, 10 543, 13 630 and so on. The jump you are hunting for is the large one: there is a modest rise from the first to the second value, then a very large jump to the third. That pattern — two comparatively low values, then a step up by a factor of roughly five — tells you two electrons are removed relatively easily from the outer shell before the third electron has to come from an inner, much more strongly held shell. Two easily removed electrons means Group 2. The element here is magnesium.
The reasoning you must write out is: "The large jump between the 2nd and 3rd ionisation energies shows the 3rd electron is removed from a shell closer to the nucleus, with less shielding, so far more energy is required. Two electrons removed before that jump means two electrons in the outermost shell, so the element is in Group 2." Notice the mark-bearing words: closer to the nucleus, less shielding, outermost shell. Candidates who simply write "the third electron is harder to remove" without the shell-and-shielding explanation lose the reasoning mark.
Explaining why Mg has a higher first ionisation energy than Al
This is the single most-set periodicity explanation item, and it trips up candidates who expect ionisation energy to rise smoothly across a period. Magnesium is [Ne]3s2; aluminium is [Ne]3s23p1. In magnesium the electron removed comes from a 3s orbital; in aluminium the electron removed comes from a 3p orbital, which is at a slightly higher energy and is slightly shielded by the 3s electrons. That extra shielding and higher orbital energy outweigh the increase of one in nuclear charge, so aluminium's outer electron is easier to remove — its first ionisation energy is lower. The full-mark answer names the sub-shell change explicitly (3s to 3p), not just "the electron is easier to remove".
Group 2: comparing the vigour of two reactions with water
Consider calcium and magnesium reacting with water. Magnesium reacts only very slowly with cold water (a few bubbles of hydrogen over days) but burns vigorously in steam:
Mg+H2O→MgO+H2
Calcium reacts steadily with cold water, giving a cloudy suspension of slightly-soluble calcium hydroxide and a stream of hydrogen:
Ca+2H2O→Ca(OH)2+H2
The examinable explanation is that reactivity increases down the group because the first and second ionisation energies both fall — the outer electrons are further from the nucleus and better shielded, so the M2+ ion forms more readily and the metal is oxidised more easily. Note the two different products: magnesium in steam gives the oxide (water is present as vapour, not liquid), whereas calcium in cold water gives the hydroxide. Getting the right product for the right conditions is a frequent discriminator. To go deeper on the redox bookkeeping behind "the metal is oxidised", revisit Acids, Redox, Electrons and Bonding.
Group 2 solubility: the two opposite trends
Students routinely muddle the two Group 2 solubility trends because they run in opposite directions. Commit both to memory as a pair:
| Compound type | Solubility trend down Group 2 | Consequence |
|---|---|---|
| Hydroxides M(OH)2 | Increases (Mg(OH)₂ almost insoluble → Ba(OH)₂ fairly soluble) | Mg(OH)₂ is a mild antacid; Ba(OH)₂ gives a strongly alkaline solution |
| Sulfates MSO4 | Decreases (MgSO₄ very soluble → BaSO₄ essentially insoluble) | BaSO₄ insolubility is the basis of the sulfate test and of the "barium meal" |
At A-Level the observation of the trend is examinable at Paper 1; the explanation in terms of the competition between lattice enthalpy and hydration enthalpy is developed quantitatively in Energetics and Electrode Potentials. The one-line version to reproduce under time pressure: as the cation gets larger down the group, both lattice enthalpy and hydration enthalpy become less exothermic, but they fall at different rates, so which effect dominates decides whether solubility rises or falls. Do not attempt to explain the trend from ionic radius alone — that is pitfall five below.
Halogen displacement: predicting the observation
You must be able to predict both the equation and the visible colour change for any halogen–halide displacement. Chlorine is a stronger oxidising agent than bromine, so it will displace bromide from solution:
Cl2+2KBr→2KCl+Br2
The observation is that a colourless solution turns orange as aqueous bromine forms. Add an organic solvent such as cyclohexane and shake, and the bromine concentrates into the upper organic layer as an orange colour; iodine under the same treatment would give a violet/purple organic layer, and chlorine a very pale green. Being able to state the colour in the organic layer — orange for bromine, violet for iodine — is often the mark that separates a top-band answer from a middling one, because the aqueous colours (yellow-orange bromine, brown iodine) are easier to confuse. The underlying trend to quote is that oxidising power decreases F2>Cl2>Br2>I2 because the atoms get larger down the group, so the incoming electron is added further from the nucleus and less strongly attracted.
Halide reducing power with concentrated sulfuric acid
The converse trend — reducing power increasing Cl−<Br−<I− — is examined through the reactions of solid sodium halides with concentrated sulfuric acid, and it produces a memorable escalation of products:
- Sodium chloride: only an acid–base reaction. Steamy white fumes of hydrogen chloride; chloride is too weak a reducing agent to reduce sulfur. NaCl+H2SO4→NaHSO4+HCl.
- Sodium bromide: HBr forms first, but bromide is a strong enough reducing agent to reduce some of the sulfuric acid, so you also see brown bromine vapour and colourless, choking sulfur dioxide (sulfur reduced from +6 to +4).
- Sodium iodide: iodide is the strongest reducing agent of the three, so reduction goes furthest — you can see violet iodine vapour, a black solid, the smell of bad eggs from hydrogen sulfide (sulfur reduced all the way from +6 to −2), plus SO2 and S as intermediate products.
The examiner reward here is tracking the oxidation state of sulfur through each stage and stating explicitly that the halide is oxidised while sulfur is reduced. A model half-equation worth reproducing for the iodide case is H2SO4+8H++8e−→H2S+4H2O. This chemistry connects forward to the reducing-power reasoning that underpins halide reactivity in Alcohols and Haloalkanes.
Exam Technique for Module 3.1
Beyond knowing the chemistry, the marks on this module are won and lost on how you phrase the answer. A few habits pay for themselves repeatedly.
- "Explain" means give the reason, not the observation. If a question says "explain why first ionisation energy increases across Period 3", a description of the trend earns nothing on its own — you must name the causes (increasing nuclear charge, same shell, similar shielding, so electrons held more strongly). Underline the command word before you start writing.
- Name all four ionisation-energy factors, then pick the dominant one. The four are nuclear charge, distance of the electron from the nucleus, shielding by inner shells, and spin-pair repulsion within a sub-shell. Almost every explanation item is answered by identifying which of these four is doing the work in the specific case, and which is being outweighed.
- State the property being used, not just the result. In descriptive chemistry, "Ba2+ forms more readily because ionisation energies are lower down the group" scores; "barium is more reactive" alone does not. The mark is for the causal link.
- Balance equations and show state symbols where solubility or gas evolution is the point. In the anion tests and the Group 2 reactions, the state symbol carries meaning — (s) for a precipitate, (g) for effervescence — and dropping it can cost a mark.
- For disproportionation, prove it. Do not just assert that a reaction is disproportionation; show the two oxidation states of the same element, one going up and one going down. For Cl2+2NaOH→NaCl+NaClO+H2O, chlorine goes from 0 to −1 (in NaCl) and from 0 to +1 (in NaClO). Writing both changes explicitly is what earns the mark.
Common-Mistake Callouts
Callout — the anion-test order is not arbitrary. Always test carbonate first, then sulfate, then halide. If you test for sulfate first with barium chloride, a carbonate impurity gives a white precipitate of barium carbonate and a false positive. Removing carbonate first (it fizzes away with dilute acid) prevents this. If you test for halide with silver nitrate without first acidifying with dilute nitric acid, a carbonate impurity gives a cream silver-carbonate precipitate that masquerades as a halide. Acidify first, every time.
Callout — "gets bigger" is not an explanation. Saying atomic radius "increases down the group because atoms get bigger" is circular and scores nothing. The explanation is that each successive element has an additional occupied electron shell, and the extra shielding from inner shells outweighs the increase in nuclear charge, so the outer electrons sit further out.
Callout — the sulfur dip does not reverse the trend. After the nitrogen-to-oxygen (or phosphorus-to-sulfur) dip, first ionisation energy resumes rising. Sulfur's first ionisation energy is lower than phosphorus's but still higher than silicon's. Answers that claim the trend "reverses" at sulfur are wrong; it merely dips before continuing upward.
Mini-FAQ
Is scandium a transition metal, and why does that matter here? Not by the strict definition — but that belongs to Module 5.3, not periodicity. Within Module 3.1 you only need the s/p/d/f block organisation. The distinction (d-block vs "true" transition metal) is developed in Transition Elements and Aromatic Chemistry.
Do I need to learn the ionisation energies as numbers? No. You need to recognise patterns — the shape of the Period 3 graph with its two dips, and the "big jump" in successive ionisation energies that reveals the group. Learn the reasoning, not a data table.
Why is beryllium the odd one out in Group 2? Beryllium is anomalously unreactive (it does not react with water even as steam) and its compounds show more covalent character because Be2+ is very small and highly polarising. At H432 you are expected to know that reactivity increases down the group starting from a low base at Be/Mg.
How does periodicity connect to the rest of the course? The controlling factors — nuclear charge, distance, shielding — reappear everywhere: in the lattice and hydration enthalpies of Energetics and Electrode Potentials, in the electrode-potential ordering that predicts halogen displacement, and in the bond-polarity arguments running through Basic Organic and Hydrocarbons. Periodicity is the reasoning engine, not a standalone topic.
Synoptic Links
The periodicity and reactivity reasoning developed here threads forward through the spec. Ionisation energy underpins lattice enthalpy calculations in energetics and electrode potentials, and the Group 2 metal-to-ion process is the standard worked example of a Born-Haber cycle. The halogen oxidising-power trend is the basis of redox-potential predictions in the same course, and the halide chemistry returns in the haloalkane substitution rates of alcohols and haloalkanes, where C-I bonds hydrolyse fastest and C-F slowest. The Period 3 structure-and-bonding story is the inorganic counterpart of the organic structure-property reasoning that runs through basic organic and the spectroscopy chapters.
Paper 3 'Unified chemistry' items deploy this module in two characteristic ways. The first is the descriptive-chemistry application: candidates are given an industrial or biological scenario (limestone neutralisation of acidic lakes with Ca(OH)₂; the use of Mg(OH)₂ as an antacid; the chlorination of drinking water and the disproportionation chemistry that controls free chlorine concentration) and asked to apply the descriptive trends. The second is the trend-justification synoptic: candidates are given physical data (a melting-point dataset, a solubility dataset, an ionisation-energy dataset) for an unfamiliar set of elements or compounds and asked to interpret it using the periodicity framework. The discriminating moves at the top band are explicit identification of which trend factor (nuclear charge, distance, shielding, lattice/hydration competition) drives the observation, and the explicit acknowledgement of the trade-off when two factors compete (the Group 2 sulfate solubility decrease versus the hydroxide increase is the canonical example).
What Examiners Reward
Top-band marks on this module cluster around precision of trend explanation and explicit use of the controlling-factor vocabulary. For first-ionisation-energy questions, examiners want the four factors named (nuclear charge, distance, shielding by inner shells, spin-pair repulsion within a sublevel) and the dominant factor identified for the specific case. For Group 2 reactivity questions, they want the link from low first ionisation energy down the group to greater willingness to form M²⁺ ions, and the consequent vigour of reaction with water or acid. For halogen oxidising-power questions, they want the link from atomic radius down the group to weaker attraction for the incoming electron, hence reduced ability to oxidise. For qualitative anion-test questions, they want the test order justified (carbonate before sulfate, so the CO₃²⁻-Ba²⁺ false positive is eliminated by acidification) and the observation described with both colour and any solubility behaviour in ammonia.
Common pitfalls cluster around six recurring mistakes. First, describing trends as "atoms get bigger down the group" without acknowledging that shielding outweighs the increasing nuclear charge — the standard exam request is to explain why, not just to state. Second, attributing the Mg-Al ionisation-energy dip to "fewer electrons" rather than to the 3s-to-3p sublevel change. Third, attributing the N-O (or P-S) dip to "the new electron entering a paired orbital" without naming the consequent spin-pair repulsion. Fourth, omitting the acidification step in the halide test (without HNO₃, a CO₃²⁻ contaminant would give a false silver-carbonate precipitate). Fifth, predicting solubility trends from atomic radius alone without referencing the lattice-versus-hydration enthalpy balance. Sixth, writing chlorine disproportionation with chlorine going to two products of the same oxidation number rather than the required +1 and -1 (the definition of disproportionation is that one species is simultaneously oxidised and reduced). Each is a one- or two-mark deduction that compounds quickly across multi-part descriptive questions.
Practical Activity Groups (PAGs)
This course anchors PAG 4 (Qualitative analysis of ions) in full. The qualitative anion analysis lesson develops the carbonate/sulfate/halide test sequence with the acidification and confirmatory ammonia solubility steps that distinguish AgCl, AgBr and AgI. The course also previews PAG 5 (Synthesis of an organic solid) indirectly by establishing the halogen and halide chemistry that later reappears in haloalkane synthesis. The Group 2 reactivity lesson sets up the observational vocabulary (effervescence, precipitate colour, solution colour) that recurs through every subsequent practical group.
Going Further
Undergraduate analogues of this material extend in two directions. First, the periodic trends generalise into Slater's rules for effective nuclear charge and then into the relativistic corrections needed to explain the heavier elements (Au's colour, Hg's liquidity at room temperature). Second, descriptive inorganic chemistry generalises into coordination chemistry (foundational for transition elements), bioinorganic chemistry (Ca²⁺ as a signalling ion, Mg²⁺ at the centre of chlorophyll) and industrial inorganic chemistry (the Solvay process, the chlor-alkali industry). Oxbridge-style interview prompts on this material include: "Why does Mg have a higher first ionisation energy than Al?" "Why does the solubility of Group 2 hydroxides increase down the group while that of the sulfates decreases?" "Predict what happens when fluorine is bubbled into water — and explain why it differs from chlorine in water."
Authorship and Sign-off
This guide was authored independently by John Haigh, paraphrasing OCR H432 Modules 3.1.1, 3.1.2 and 3.1.3 as descriptive use. No verbatim spec text, mark-scheme phrasing, examiner-report quotation, or past-paper question reference appears. The worked examples are original.
Start at the Periodicity, Group 2 and Halogens course and work through every lesson in sequence. Once the trend logic and the descriptive chemistry of Group 2 and the halogens are automatic, every later H432 module becomes a structural extension of the same ideas — and the descriptive-inorganic questions become recognition rather than recall.