The Periodic Table and s, p, d Blocks

Spec Mapping — OCR H432 Module 3.1.1 — Periodicity, content statements covering the arrangement of the modern Periodic Table by atomic (proton) number, the definitions of period, group and block, classification of elements into s-, p-, d- and f-blocks according to the highest-energy sub-shell occupied, and the link between outer electron configuration and chemical similarity within a group (refer to the official OCR H432 specification document for exact wording). This lesson is the gateway to the whole of Module 3.1; every later argument about reactivity trends, ionisation energy and Group 2/Group 7 chemistry depends on the block framework established here.

The Periodic Table is the single most powerful organising idea in chemistry. Roughly 60 % of the marks in OCR H432 Paper 1 ("Periodic table, elements and physical chemistry") rely on you reading the table correctly — identifying the block, predicting the type of ion an element will form, ranking elements by ionisation energy, and inferring oxidation-state patterns. In this lesson we re-derive the modern Periodic Table from the quantum-mechanical picture established in Module 2.2.1 (electron configurations), so that the table is no longer a wall poster to memorise but a map of electron filling. The conceptual heart of the lesson is the realisation that the column an element occupies is fixed by the configuration of its outer sub-shell, and that the outer sub-shell drives every chemical property that varies across the Periodic Table — from melting point to reaction with water to the colour of the resulting solution.

Key Definitions:

• Period — a horizontal row of the Periodic Table; equals the principal quantum number $n$n of the highest occupied shell.
• Group — a vertical column; elements in the same group share the same outer-shell electron pattern.
• Block — a region of the table named after the sub-shell into which the highest-energy electron of the element is placed (s, p, d or f).
• Periodicity — the repeating pattern of chemical and physical properties seen as you move across successive periods.

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Arrangement by atomic number

The modern Periodic Table is arranged in order of increasing atomic (proton) number ($Z$Z), not atomic mass. The shift from mass-ordering (Mendeleev's 1869 table) to proton-number ordering was made possible by Henry Moseley's 1913 X-ray emission experiments, which showed that the frequencies of characteristic X-rays scale linearly with $\sqrt{Z}$Z​:

$\sqrt{\nu} \propto (Z - \sigma)$ν​∝(Z−σ)

This Moseley's law gave each element a unique integer "atomic number" and resolved the anomalies that had puzzled Mendeleev. Two pairs in particular sit "out of order" by atomic mass:

Pair	Order by mass	Order by $Z$Z	Why correct
Ar (39.95) / K (39.10)	K, Ar (mass)	Ar (18), K (19)	Chemical properties match Ar with noble gases, K with alkali metals
Te (127.6) / I (126.9)	I, Te (mass)	Te (52), I (53)	Te is a metalloid (Group 16); I a halogen (Group 17)
Co (58.93) / Ni (58.69)	Ni, Co (mass)	Co (27), Ni (28)	Their chemistries belong in those positions

The atomic mass pairs invert because of natural isotope abundance — Ar has unusually heavy isotopes (mainly Ar-40) and K has unusually light ones (mainly K-39). Ordering by $Z$Z eliminates the anomaly because $Z$Z is the proton count and is identical for every atom of an element.

flowchart LR
    A[Mendeleev 1869<br/>ordered by atomic mass] --> B[Moseley 1913<br/>X-ray frequencies show integer Z]
    B --> C[Modern Periodic Table<br/>ordered by atomic number]
    C --> D[Periods: principal shell n]
    C --> E[Groups: outer e- pattern]
    C --> F[Blocks: highest sub-shell occupied]

Periods are numbered 1 to 7 from the top; groups are numbered 1 to 18 by the modern IUPAC system, although OCR retains the older "Group 2, Group 7" labels for s- and p-block columns. Period number equals the principal quantum number ($n$n) of the highest occupied shell — Period 3 elements have outer electrons with $n = 3$n=3; Period 4 with $n = 4$n=4; and so on.

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Periodicity — the repeating pattern

Periodicity is the repeating pattern of physical and chemical properties seen as you traverse successive periods. The reason the pattern repeats is that each new period begins with an electron entering the next-higher s sub-shell (so a fresh outer shell is started), and the chemistry is dictated by that outer shell. Period 2 (Li → Ne) and Period 3 (Na → Ar) both run from "one outer s electron" through "three outer p electrons" to "full outer s²p⁶" — so the chemistries echo each other.

Examples of periodic properties you will study in this module:

• Atomic radius — decreases across a period, increases down a group (next lesson).
• First ionisation energy — generally increases across a period (with two characteristic dips at Group 13 and Group 16), decreases down a group.
• Electronegativity — increases across a period (F is the most electronegative), decreases down a group.
• Melting point — peaks at Group 14 (giant covalent: C, Si) and falls sharply for Groups 15–18 (simple molecular).
• Type of structure and bonding — giant metallic → giant covalent → simple molecular.
• Oxide acidity — basic on the left (Na₂O), amphoteric in the middle (Al₂O₃), acidic on the right (SO₃, Cl₂O₇).

The periodic table makes chemistry predictable. Once you have established the trend on one row, you can predict the corresponding behaviour on the next — and this is exactly what Mendeleev did in his 1869 paper.

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Classification by electron configuration: s, p, d, f blocks

The Periodic Table is partitioned into four blocks, each named after the type of sub-shell into which the highest-energy electron of the element is placed:

s-block: outermost electron in s sub-shell (ns¹ or ns²) p-block: outermost electron in p sub-shell (ns² np¹⁻⁶) d-block: outermost electron entering (n−1)d f-block: outermost electron entering (n−2)f

Block	Groups	Outer configuration	Examples
s	1, 2 (and He by configuration)	$ns^1$ns1 or $ns^2$ns2	Na [Ne] 3s¹, Mg [Ne] 3s²
p	13 – 18	$ns^2\, np^{1-6}$ns2np1−6	Cl [Ne] 3s² 3p⁵, Ar [Ne] 3s² 3p⁶
d	Transition metals (Sc → Zn, Y → Cd, La → Hg)	$(n-1)d^x\, ns^2$(n−1)dxns2	Fe [Ar] 3d⁶ 4s², Cu [Ar] 3d¹⁰ 4s¹
f	Lanthanides + actinides	$(n-2)f^x\, (n-1)d^{0\text{ or }1}\, ns^2$(n−2)fx(n−1)d0 or 1ns2	U [Rn] 5f³ 6d¹ 7s²

Helium is a special case: its electron configuration is 1s² and it sits formally in the s-block by configuration, but chemically it is grouped with the noble gases (Group 18) because its single outer shell is full. OCR expects you to recognise this anomaly and treat helium as a noble gas in chemical contexts.

Why the block rule works

In Module 2.2.1 (electron configurations) you learned the Aufbau filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d. The Periodic Table is, in effect, a map of this filling order laid out so that each new row starts a new s sub-shell and the position across the row corresponds to the next electron added. Sc (Z = 21) is the first element to receive a d electron, so it starts the d-block. Once you internalise this — the table is the filling order — block classification becomes automatic.

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Why group members behave similarly

The chemistry of an atom is dominated by its valence electrons — the electrons in the outermost (highest-$n$n) shell, plus, for transition metals, the partly-filled (n−1)d. Atoms in the same group share the same valence-electron count, so they form the same type of ion and undergo the same characteristic reactions.

Element	Z	Configuration	Outer electrons
Be	4	[He] 2s²	2 (in s)
Mg	12	[Ne] 3s²	2 (in s)
Ca	20	[Ar] 4s²	2 (in s)
Sr	38	[Kr] 5s²	2 (in s)
Ba	56	[Xe] 6s²	2 (in s)

All Group 2 elements have two electrons in their outer s sub-shell. They all lose these two electrons during reaction to form 2+ ions with the configuration of the preceding noble gas, and they all exhibit the same reactions (with water, with oxygen, with halogens). The same principle works for Group 17 (the halogens — gain one electron to form X⁻), for Group 1 (lose one electron to form M⁺), and for Group 18 (noble gases — already full outer shell, very low reactivity).

The strength of this argument is that it explains why the periodic table works — without it, the table is a list to memorise; with it, the table is a deduction tool.

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Mendeleev's predictive table

Dmitri Mendeleev (1869) arranged the then-known 63 elements primarily by atomic mass into a table he called the "periodic system". Two innovations gave his table predictive power:

1. He left gaps where he believed undiscovered elements should sit, and predicted the properties of those missing elements from the trends of their neighbours.
2. He placed Te before I despite the mass ordering being I before Te, on the basis that the chemistry "felt right" with Te in Group 16 (oxygen family) and I in Group 17 (halogen family).

His famous prediction of "eka-silicon" — an undiscovered Group 14 element with predicted atomic mass ~72, density ~5.5 g cm⁻³, oxide formula MO₂ — was vindicated by the 1886 discovery of germanium by Clemens Winkler, with measured properties accurate to within a few percent of Mendeleev's forecasts (mass 72.6, density 5.32, oxide GeO₂). Other predictions — eka-aluminium (gallium, 1875) and eka-boron (scandium, 1879) — also bore out. These successes are what entrenched the periodic system as a predictive tool, not merely an organising one.

The modern table differs from Mendeleev's in two structural respects:

1. Elements are ordered by atomic number (Moseley, 1913), not atomic mass — eliminating the Ar/K, Co/Ni and Te/I anomalies.
2. A separate d-block (transition metals) was added after Mendeleev's death; the f-block (lanthanides and actinides) was added later still.

The deeper reason periodicity works is quantum-mechanical (sub-shells with characteristic capacities of 2, 6, 10, 14 electrons) — but Mendeleev had no access to that theory in 1869, and his table was nonetheless astonishingly correct. This is one of the great empirical triumphs of nineteenth-century chemistry.

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Worked examples

Example 1 — block classification by configuration

Q: Classify each of the following elements by block and state the outer electron configuration: (a) sulfur (Z = 16); (b) vanadium (Z = 23); (c) strontium (Z = 38); (d) gallium (Z = 31); (e) cerium (Z = 58, lanthanide).

Answer:

(a) S: 1s² 2s² 2p⁶ 3s² 3p⁴ → outer sub-shell is 3p → p-block; outer 3s² 3p⁴. (b) V: [Ar] 3d³ 4s² → highest occupied sub-shell is 3d → d-block (transition metal); outer 3d³ 4s². (c) Sr: [Kr] 5s² → outer 5s → s-block (Group 2); outer 5s². (d) Ga: [Ar] 3d¹⁰ 4s² 4p¹ → outer 4p → p-block (Group 13); outer 4s² 4p¹. Note the 3d¹⁰ is full and chemically inert at the AS level. (e) Ce: [Xe] 4f¹ 5d¹ 6s² → highest occupied sub-shell is 4f → f-block (lanthanide); outer 4f¹ 5d¹ 6s².

Example 2 — predict an unknown

Q: Element X has the configuration [Kr] 4d¹⁰ 5s² 5p⁵. Identify (i) the block; (ii) the group; (iii) the period; (iv) the most likely ion it forms.

Answer: (i) p-block (outer in 5p). (ii) Group 17 (halogen) — five 5p electrons make seven outer electrons in total. (iii) Period 5 (highest $n$n = 5). (iv) X⁻, the −1 ion, by gaining one electron to complete the 5p sub-shell. The element is iodine (Z = 53), and you would predict iodine to behave like Br, Cl and F — forming hydrogen halides, displacing less reactive halides from solution, etc.

Example 3 — anomaly check

Q: A student claims Ca is a d-block element because it lies in Period 4 and "the transition metals are in Period 4". Explain why the student is wrong.

Answer: Calcium's electron configuration is [Ar] 4s² — its highest-energy electron occupies the 4s sub-shell, not 3d. By the block-classification rule (block = highest sub-shell occupied), Ca is therefore an s-block element in Group 2. The transition-metal d-block proper begins one place to the right at scandium (Sc, [Ar] 3d¹ 4s²). Ca's chemistry — losing two electrons to form Ca²⁺, reacting with water to give Ca(OH)₂ — is classic Group 2 behaviour, not transition-metal behaviour. (Transition metals also typically show variable oxidation states and coloured ions, neither of which Ca does.)

Example 4 — diagonal relationships (extension)

Q: Lithium (Group 1, Period 2) shows several similarities to magnesium (Group 2, Period 3). Account for this in terms of position and electron arrangement.

Answer: Lithium and magnesium lie on a diagonal in the Periodic Table. Although they are in different groups, they have similar charge density (charge / ionic radius): Li⁺ is small (76 pm) and singly charged; Mg²⁺ is slightly larger (72 pm) but doubly charged, giving a similar polarising power. Diagonal pairs (Li/Mg, Be/Al, B/Si) therefore share several chemical features — for instance, both Li and Mg form simple ionic nitrides on burning in nitrogen (Li₃N and Mg₃N₂), whereas Na (Li's group-mate) does not. Diagonal relationships are not assessed for recall by OCR but are a standard Oxbridge-interview talking point and show how polishing periodic patterns illuminates exceptions.

Example 5 — Mendeleev's missing element

Q: Mendeleev's 1871 table left a gap below silicon in Group 14 (his old "Group IV"). He named the missing element "eka-silicon" and predicted properties from its neighbours Si, Sn, Zn, As. State two predicted properties and one observed property of the element that filled the gap.

Answer: Two predicted properties: (i) atomic mass ~72; (ii) oxide formula MO₂ with high melting point; one observed property of the gap-filler germanium (discovered 1886): atomic mass 72.6, oxide GeO₂ melting at 1115 °C. Mendeleev's prediction matched experimental values to within a few percent, which contemporaries treated as decisive evidence in favour of the periodic system.

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Synoptic Links

Synoptic Links — Connects to:

• ocr-alevel-chemistry-acids-redox-bonding / electron-configurations-and-orbitals (the Aufbau filling order is what generates the s/p/d/f blocks; understanding orbital energies is prerequisite).
• ocr-alevel-chemistry-periodicity-groups / periodic-trends-atomic-radius-ionisation-energy (block position predicts ionisation-energy trends — the very next lesson).
• ocr-alevel-chemistry-atoms-moles / mass-spectrometry (Moseley's law uses X-ray frequencies to assign Z; mass spectrometry assigns atomic mass — the two pieces of data together fully characterise an isotope).
• ocr-alevel-chemistry-transition-aromatic / d-block-properties (Year 13 transition-metal chemistry is the natural continuation of the d-block introduction here).

Practical Activity Group anchor: No direct PAG anchor for the block-classification topic itself, but PAG 4 (qualitative analysis of cations and anions) implicitly relies on block knowledge — Group 2 cations behave differently from Group 1, transition-metal cations give characteristic colours, and halide anions behave by group. The qualitative-analysis lesson at the end of this course revisits this point.

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Specimen question modelled on the OCR H432 paper format

Question (6 marks): An unknown element X has the electron configuration [Kr] 4d¹⁰ 5s² 5p³. (a) Identify the block, group and period of X. (3 marks) (b) Predict, with reasoning, two chemical properties of X based on its position. (3 marks)

AO breakdown

Mark	AO	Awarded for
1	AO1	Block: p-block (highest sub-shell occupied is 5p)
2	AO1	Group: 15 (or Group 5) — five outer electrons (5s² 5p³)
3	AO1	Period: 5 (highest $n$n = 5)
4	AO2	Property 1: forms a 3− ion (X³⁻) by gaining three electrons to complete 5p, or forms covalent compounds with H giving XH₃
5	AO2	Property 2: oxide formula X₂O₅ or X₂O₃ (Group 15 maximum oxidation state +5; lower +3)
6	AO3	Reasoning links block/group position to outer-electron count and to the periodicity of oxidation states

AO split: AO1 = 3, AO2 = 2, AO3 = 1.

Mid-band response

(a) The highest occupied sub-shell is 5p, so X is in the p-block. It has five outer electrons (5s² 5p³), so it is in Group 15 (using IUPAC numbering, Group 5 in the older notation). The principal quantum number of the outer shell is 5, so X is in Period 5.

(b) Two predicted properties:

1. X is likely to gain three electrons to form a 3− anion (X³⁻), or to form three covalent bonds in compounds with hydrogen, giving the formula XH₃ (compare NH₃, PH₃).
2. X can show oxidation states of −3, +3 and +5 — Group 15 elements typically have an oxide of formula X₂O₅ (highest oxidation state) and X₂O₃ (intermediate). The element X is in fact antimony (Sb, Z = 51).

Examiner commentary: M1 (AO1) block; M1 (AO1) group; M1 (AO1) period; M1 (AO2) property 1 ion/covalency; M1 (AO2) property 2 oxide formula. The candidate skirts the full reasoning mark by naming Sb at the end but not explicitly linking "Group 15 → five outer electrons → variable oxidation states ±3 / +5". Around 5/6.

Top-band response

(a) The highest-energy sub-shell occupied is 5p (the configuration [Kr] 4d¹⁰ 5s² 5p³ has its last electron entering 5p), so X belongs to the p-block. The outer-shell count is 2 + 3 = 5 valence electrons, placing X in Group 15 of the modern Periodic Table (older "Group V"). The principal quantum number of the outer shell is $n = 5$n=5, hence Period 5.

(b) Two predictions that follow from this position:

1. Bonding behaviour. With five valence electrons, X needs three more to complete its octet, so it is likely to form three covalent bonds. By analogy with nitrogen and phosphorus, X will form a hydride of formula XH₃ — pyramidal in shape (VSEPR: three bond pairs + one lone pair on the central atom). The 3− anion X³⁻ is also possible in ionic contexts (e.g. in nitrides), although for the heavier Group 15 elements covalent bonding dominates.
2. Oxidation states and oxides. Group 15 elements show variable oxidation states, characteristically −3 (in hydrides), +3 (in trihalides such as XCl₃), and +5 (in pentahalides such as XCl₅, and in oxides such as X₂O₅). The maximum oxidation state of +5 reflects the involvement of all five valence electrons. Going down Group 15, the +3 oxidation state becomes progressively more stable relative to +5 (the "inert pair effect" — the s² pair becomes harder to involve in bonding as the s sub-shell drops further below the p in energy at high $Z$Z). For X at $Z = 51$Z=51 (antimony), Sb(III) compounds are slightly more stable than Sb(V); for the next element down (Bi, $Z = 83$Z=83), Bi(III) dominates almost entirely.

Examiner commentary: Full 6/6 with comfort. The discriminators that lift this answer to A*: (i) explicit naming of the inert pair effect as a periodic trend, (ii) explicit VSEPR reasoning for the XH₃ geometry, (iii) explicit synoptic link from Group 15 to oxidation-state stability ordering. The candidate has gone well beyond rote prediction into Year-13-level reasoning.

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Common errors and mark-loss patterns

Pedagogical observations — not fabricated statistics:

• A frequent pitfall is classifying Ca as d-block because it sits in Period 4 alongside the transition metals. Ca is [Ar] 4s² — outermost electron in 4s — so it is s-block, Group 2.
• Many candidates write Fe as [Ar] 4s² 3d⁶ in exam answers. OCR strictly prefers the order [Ar] 3d⁶ 4s² because 3d is the higher-shell-number sub-shell and 4s electrons are removed first on ionisation; writing 4s before 3d obscures the ionisation order.
• A typical pitfall is forgetting the He anomaly: helium is 1s² (s-block by configuration), but is chemically a noble gas and sits in Group 18 (Group 0 in older notation).
• Candidates sometimes confuse "period" and "group" in vocabulary — period is the row, group is the column. The mnemonic "Group = Going down, Period = horizontal" helps if you remember it.
• Many lose marks by claiming the modern table is "ordered by atomic mass" — it is not. Moseley's law established atomic-number ordering in 1913.
• A subtle error is mis-assigning the block for lanthanides and actinides: cerium ([Xe] 4f¹ 5d¹ 6s²) is f-block, not d-block, even though it has a 5d¹ electron. The rule "highest sub-shell" means highest in terms of being the latest to fill, which is 4f, not 5d, by the Aufbau order.
• Some candidates omit "(g)" state symbols in ionisation-energy contexts, or describe Mendeleev's table as "the modern table" — both are easy ways to lose a mark.

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Going further

The block structure of the Periodic Table emerges from solutions of the Schrödinger equation for hydrogen-like atoms: the angular-momentum quantum number $\ell$ℓ takes values $0, 1, 2, 3$0,1,2,3, which we label s, p, d, f, and each value of $\ell$ℓ admits $2\ell + 1$2ℓ+1 orbitals containing up to $2(2\ell + 1)$2(2ℓ+1) electrons. The s-block is therefore 2 columns wide, the p-block 6, the d-block 10, the f-block 14 — all because the Schrödinger equation has those degeneracies. The Madelung rule (filling in order of increasing $n + \ell$n+ℓ, then increasing $n$n) gives the Aufbau order and the block layout. Beyond AS, you may meet density-functional theory (DFT) (Walter Kohn, Nobel Prize 1998), which can reproduce the entire periodic table from first principles with no empirical input. An Oxbridge-interview-style prompt: "Why is the f-block separated and printed below the main table?" — the answer is purely typographic (the table would be 32 columns wide if drawn in full), and you should be able to explain that the f-block proper sits between Groups 3 and 4 in Periods 6 and 7. Recommended reading: Eric Scerri's The Periodic Table: Its Story and Its Significance (Oxford) — a full intellectual history; Atkins, Inorganic Chemistry (Atkins, Overton, Rourke et al.) for the quantum-mechanical foundation.

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A-Level misconceptions

The errors that distinguish A from A*:

1. The Periodic Table is ordered by atomic number, not atomic mass. This was Moseley's contribution in 1913, post-Mendeleev. Mendeleev's table was mass-ordered with a chemical-intuition over-ride (Te before I).
2. Block ≠ group. The s-block contains Groups 1 and 2 (two columns); the p-block Groups 13–18 (six columns); the d-block ten transition-metal columns; the f-block fourteen lanthanide/actinide columns.
3. He is in the s-block by configuration but the noble-gas column chemically. Always treat helium as a noble gas in chemical contexts.
4. Period number = principal quantum number of the outer shell, not the row index "started counting from 1". Period 1 has $n = 1$n=1; Period 4 has $n = 4$n=4.
5. Transition metals fill 4s before 3d in neutral atoms, but lose 4s first on ionisation. Fe is [Ar] 3d⁶ 4s²; Fe²⁺ is [Ar] 3d⁶ (4s gone first).
6. Cr and Cu show anomalous configurations: Cr is [Ar] 3d⁵ 4s¹, Cu is [Ar] 3d¹⁰ 4s¹ (not 3d⁴ 4s² and 3d⁹ 4s²). Half-filled and filled d sub-shells offer extra stability.
7. Periodicity is a horizontal phenomenon (across periods), not a vertical one (down groups). Groups give similar chemistry; periods give patterns of variation.
8. Group number tells you the number of outer electrons (s + p), not the row number or the period. Sulfur is Group 16: it has six outer electrons (3s² 3p⁴), and forms S²⁻ by gaining two to complete the octet.

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Summary

The modern Periodic Table arranges elements in order of increasing atomic number (Moseley's rule) into seven periods (rows) and eighteen groups (columns). Periodicity — the repeating pattern of physical and chemical properties — emerges because each new period starts a new outer shell and the chemistry is dictated by that outer shell. Elements are partitioned into four blocks (s, p, d, f) according to the type of sub-shell housing the highest-energy electron, and the dimensions of those blocks (2, 6, 10, 14 columns) follow from the quantum-mechanical capacities of s, p, d and f orbitals. Group members share an outer-electron count and therefore form similar ions and undergo similar reactions. Mendeleev's 1869 system, predictive of germanium and gallium, was vindicated by both empirical discovery and by Moseley's later atomic-number ordering and the quantum-mechanical sub-shell theory. The block framework established here underpins every subsequent lesson in OCR H432 Module 3.1, from ionisation-energy trends to Group 2 reactivity to halogen displacement.

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Reference: OCR A-Level Chemistry A (H432) Module 3.1.1 (a)–(c) (refer to the official OCR H432 specification document for exact wording).