Periodicity and Trends Across Period 3

When Mendeleev first arranged the known elements into the periodic table in 1869, he did so on the basis of repeating chemical behaviour — ordering by atomic mass and leaving gaps where his pattern demanded an undiscovered element. The modern table, ordered by atomic number and partitioned by sub-shell filling, makes the underlying logic explicit: the chemistry of an element is a direct consequence of its electron configuration. Period 3 — sodium through argon — is the single most-examined row at A-Level, because it spans the full structural sweep from reactive metal to giant-covalent network to molecular non-metal to monatomic noble gas across only eight elements. This lesson previews the five quantitative trends you must master: atomic radius, first ionisation energy (IE), electronegativity, melting point, and the acid-base behaviour of the Period 3 oxides. Each trend is grounded in two competing factors — nuclear charge and shielding — and the anomalies (the Mg → Al and P → S IE drops; the dramatic melting-point fall at phosphorus) are exactly where the highest-tariff marks are earned.

Spec mapping (AQA 7405): This lesson anchors §3.2.1 (periodicity — classification of elements in s, p, d blocks; properties of Period 3 elements Na to Ar; the structures and melting points of the Period 3 elements; the trends in atomic radius, first ionisation energy and electronegativity). It builds directly on §3.1.1.3 (the patterns in successive ionisation energies, covered in atomic-structure L3 of this course series), and extends into §3.2.3 (Period 3 oxides — covered in detail in L3 of the present course). It is cross-referenced by L1 (Group 2 trends) and L2 (Group 7 trends) of this course. Refer to the official AQA specification document for exact wording.

Assessment objectives: AO1 recall items include the definitions of periodicity, first ionisation energy, atomic radius, and electronegativity. AO2 features dominate this lesson — sketching and explaining the trend in IE across Period 3, predicting structural changes and their effect on melting points, and writing balanced equations for oxide formation. AO3 reasoning appears in the high-tariff anomaly questions: rationalising the Mg → Al IE drop (sub-shell energy difference), the P → S IE drop (3p electron pairing), and predicting trends in oxide structure and acid-base behaviour from the bonding-character continuum.

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What is Periodicity?

Periodicity is the repeating pattern of physical and chemical properties across periods of the Periodic Table. Period 3 (Na to Ar) is the canonical worked example at A-Level because every property you study — atomic radius, ionisation energy, electronegativity, melting point, oxide structure, and oxide acidity — shows a sharp, well-defined trend across just eight elements. Period 2 (Li to Ne) shows the same patterns but with second-row anomalies (e.g. unusually small atomic radii) that obscure the underlying physics; Period 3 is where the textbook trends are cleanest.

The two underlying factors are nuclear charge (number of protons, Z) and shielding (the screening of outer electrons from the nucleus by inner electrons). Across a period the nuclear charge increases by one per element, while the shielding contribution from inner shells (1s²2s²2p⁶ in Period 3) remains essentially constant. The net effective nuclear charge experienced by the outer electrons (Z_eff) therefore rises, with predictable consequences for radius, IE, and electronegativity.

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Atomic Radius across Period 3

Atomic radius decreases across Period 3 from sodium to chlorine.

Element	Na	Mg	Al	Si	P	S	Cl	Ar
Atomic radius / pm	186	160	143	117	110	104	99	71
Protons (Z)	11	12	13	14	15	16	17	18
Electron config	[Ne]3s¹	[Ne]3s²	[Ne]3s²3p¹	[Ne]3s²3p²	[Ne]3s²3p³	[Ne]3s²3p⁴	[Ne]3s²3p⁵	[Ne]3s²3p⁶

Explanation

• Across the period, each successive element has one more proton in the nucleus and one more electron in the same principal energy level (n = 3).
• The additional electron enters the same shell, so there is no significant increase in shielding (the inner-shell configuration 1s²2s²2p⁶ is constant across the row).
• The increased nuclear charge pulls the outer electrons closer to the nucleus, reducing the atomic radius.
• Argon's listed radius (71 pm) is a van der Waals radius, not a covalent or metallic radius — direct comparison with the bonded radii of Na to Cl is therefore misleading. Examiners normally compare Na to Cl only.

Exam Tip: When asked to explain the trend in atomic radius, always refer to (1) increasing nuclear charge, (2) approximately constant shielding (same inner-shell electrons), and (3) greater attraction for outer electrons. Never write "more protons attract more electrons" without specifying that shielding remains roughly constant — examiners deduct marks for omitting the shielding clause.

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First Ionisation Energy across Period 3

The first ionisation energy (IE₁) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous unipositive ions:

X(g) → X⁺(g) + e⁻

Element	Na	Mg	Al	Si	P	S	Cl	Ar
1st IE / kJ mol⁻¹	496	738	578	786	1012	1000	1251	1521

General Trend

First ionisation energy generally increases across Period 3. The reasons mirror those for atomic radius:

• Increasing nuclear charge across the period (Z rises from 11 to 18).
• Negligible change in shielding (each new electron enters the same n = 3 shell; the inner 1s²2s²2p⁶ remains constant).
• Outer electrons are held more tightly, so more energy is required to remove them.

Anomaly 1: The Mg → Al Drop (738 → 578 kJ mol⁻¹)

This is the first of two carefully-examined departures from the upward trend.

• Magnesium: outer electron is in the 3s sub-shell (configuration [Ne]3s²). The 3s orbital lies closer to the nucleus and is more penetrating.
• Aluminium: outer electron is in the 3p sub-shell (configuration [Ne]3s²3p¹). The 3p orbital is slightly higher in energy than 3s; the 3s² pair provides some additional shielding for the 3p electron.
• The 3p electron in Al is therefore easier to remove than the 3s electron in Mg, and IE₁(Al) < IE₁(Mg) despite Al having one more proton.

This pattern (IE drop on entering a new sub-shell) is general: it also appears between Be → B (Period 2) and Ca → Ga (Period 4).

Anomaly 2: The P → S Drop (1012 → 1000 kJ mol⁻¹)

The second anomaly is smaller but examined just as often.

• Phosphorus: 3p³ — each of the three 3p orbitals (3p_x, 3p_y, 3p_z) contains exactly one electron, in line with Hund's rule. There is no pairing of electrons in any 3p orbital.
• Sulfur: 3p⁴ — one of the 3p orbitals now contains a pair of electrons.
• The paired electrons in sulfur experience spin-pair repulsion (mutual electrostatic repulsion when forced into the same orbital), which raises their energy slightly and makes one easier to remove.
• Therefore IE₁(S) < IE₁(P), again despite S having one more proton.

The same pattern recurs between N → O (Period 2) and As → Se (Period 4): a half-filled p sub-shell is energetically favoured because spin-pair repulsion is minimised, so the next element loses one electron unusually easily.

Exam Tip: These two dips are examined frequently. For the Mg → Al dip, emphasise sub-shell energy levels (3p > 3s) and slight extra shielding from the 3s² pair. For the P → S dip, emphasise spin-pair repulsion in the doubly-occupied 3p orbital. Always quote the electron configurations explicitly — examiners reward this.

Worked Example: Comparing IE across Periods

Q: Explain why the first ionisation energy of sodium (496 kJ mol⁻¹) is much lower than that of neon (2081 kJ mol⁻¹).

A: Sodium's outer electron occupies the 3s sub-shell, which is much further from the nucleus than neon's outer 2p electrons. Sodium has an additional inner shell (the n = 2 shell, with eight electrons) providing substantial extra shielding. Although sodium has one more proton (Z = 11 vs Z = 10 for Ne), the combined effect of greater radial distance and significantly increased shielding far outweighs the small extra nuclear charge, so sodium's first ionisation energy is much lower. This drop in IE on starting a new shell defines the boundary between periods and is the most dramatic IE change anywhere in the table.

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Electronegativity across Period 3

Electronegativity (χ) is the ability of an atom to attract the bonding pair of electrons in a covalent bond towards itself. The Pauling scale is conventional in A-Level work, with fluorine assigned the maximum value of 3.98.

Element	Na	Mg	Al	Si	P	S	Cl
Electronegativity (Pauling)	0.93	1.31	1.61	1.90	2.19	2.58	3.16

Trend

Electronegativity increases smoothly across Period 3 from sodium to chlorine.

• Increasing nuclear charge with similar shielding means the nucleus attracts bonding electrons more strongly.
• Decreasing atomic radius means bonding electrons in a covalent bond sit closer to the nucleus.
• Argon is excluded because it does not normally form bonds — assigning an electronegativity is therefore not chemically meaningful.

The increase is fairly linear across Period 3, in contrast to the IE trend which has two anomalies. This is because electronegativity is an averaged property over many bonding environments — the small spin-pair-repulsion effects that show up in IE are averaged out.

Common Misconception: Electronegativity is not the same as electron affinity. Electronegativity is a relative, dimensionless scale describing an atom's bonding behaviour in a molecule; electron affinity is the enthalpy change (kJ mol⁻¹) when a gaseous atom gains an electron to form a gaseous anion. The two correlate but measure different things.

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Melting Points across Period 3

Melting points across Period 3 show a distinctive zig-zag pattern that reflects discrete changes in bonding and structure, not a smooth electronic trend. This is the most-mistaught topic at A-Level — the explanation must always start with structure.

Element	Na	Mg	Al	Si	P₄	S₈	Cl₂	Ar
Structure	Giant metallic	Giant metallic	Giant metallic	Giant covalent	Simple molecular	Simple molecular	Simple molecular	Monatomic
Melting point / °C	98	650	660	1410	44	113	−101	−189

Metals: Na, Mg, Al

Melting point increases sharply Na → Mg → Al (98 → 650 → 660 °C) because:

1. Number of delocalised electrons per atom rises: Na donates 1, Mg donates 2, Al donates 3 — strengthening the sea of delocalised electrons.
2. Cation charge rises: Na⁺, Mg²⁺, Al³⁺ — increasing the electrostatic attraction between cations and delocalised electrons.
3. Ionic radius of the cation decreases (higher charge/radius ratio), allowing closer packing.
4. The net effect is much stronger metallic bonding. Mg and Al have very similar melting points (650 vs 660 °C) because the increase in cation charge from +2 to +3 is partly offset by changes in crystal structure (Mg is hexagonal close-packed, Al face-centred cubic).

Giant Covalent: Si (1410 °C)

Silicon has by far the highest melting point in Period 3. It crystallises in a diamond-cubic structure: each Si atom is covalently bonded to four neighbours in a tetrahedral arrangement, with bond lengths of 235 pm. To melt silicon, you must break a substantial fraction of these strong covalent bonds (Si–Si bond energy 226 kJ mol⁻¹), which requires very high temperature. This single structural change — from metallic to giant covalent — accounts for almost the entire vertical spread in the Period 3 melting-point graph.

Simple Molecular: P₄, S₈, Cl₂

These elements exist as discrete molecules in the solid state, held together by weak London (dispersion) forces between molecules. Melting depends only on overcoming these intermolecular forces — the covalent bonds within each molecule remain intact. London-force strength scales with the number of electrons per molecule (more electrons → larger instantaneous dipoles → stronger forces):

• S₈ has 128 electrons per molecule → strongest London forces → highest mp of the three (113 °C).
• P₄ has 60 electrons → intermediate London forces → mp 44 °C.
• Cl₂ has 34 electrons → weakest London forces → mp −101 °C.

The drop from Si (1410 °C) to P₄ (44 °C) is the largest single step in the Period 3 melting-point pattern and is the canonical A-Level question.

Monatomic: Ar (−189 °C)

Argon exists as discrete monatomic atoms with no chemical bonding. Only very weak London forces between isolated atoms (18 electrons each) hold the solid together, giving the lowest melting point in Period 3.

Exam Tip: The single most important step in answering any melting-point question is to identify the structure type first (metallic, giant covalent, simple molecular, monatomic) and then discuss the forces broken on melting. A common mistake is trying to explain the trend using only one factor (e.g. nuclear charge) when the real explanation requires discussing the change of bonding type. Marks are routinely awarded for explicitly naming the structure.

Worked Example: Si vs P

Q: Explain why silicon has a much higher melting point than phosphorus.

A: Silicon has a giant covalent structure in which each atom is covalently bonded to four neighbours in a tetrahedral lattice. To melt silicon, many strong covalent bonds (Si–Si, 226 kJ mol⁻¹) must be broken, requiring very high temperatures (1410 °C). Phosphorus exists as discrete P₄ molecules in which atoms are covalently bonded internally but only weak London (dispersion) forces operate between molecules. Melting phosphorus requires only that the weak intermolecular forces be overcome — the strong P–P covalent bonds within each tetrahedron remain intact. Hence the dramatic drop in melting point from Si to P₄.

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Period 3 Oxides: Structure and Acid-Base Behaviour

The oxides of Period 3 elements show a parallel structural progression (giant ionic → amphoteric → giant covalent → molecular covalent) that maps onto a clear acid-base trend (basic → amphoteric → acidic). This is the direct extension into L3 of the present course (Period 3 oxides — full treatment) and is anchored in §3.2.3 of the AQA specification. A concise overview is given here to make the periodicity story complete.

Oxide	Na₂O	MgO	Al₂O₃	SiO₂	P₄O₁₀	SO₃	Cl₂O₇
Bonding	Ionic	Ionic	Ionic with covalent character	Giant covalent	Molecular covalent	Molecular covalent	Molecular covalent
Structure	Giant ionic	Giant ionic	Giant ionic (polarised)	Giant covalent	Molecular	Molecular	Molecular
Mp / °C	1132	2852	2072	1610	340	17	−9
pH in water	13–14	9–10	insoluble	insoluble	1–2	0–1	0–1
Acid-base	Strongly basic	Basic	Amphoteric	Weakly acidic	Acidic	Strongly acidic	Strongly acidic

Key Equations

• Na₂O(s) + H₂O(l) → 2 NaOH(aq) — strongly basic.
• MgO(s) + H₂O(l) → Mg(OH)₂(aq) — weakly basic; slightly soluble.
• Al₂O₃(s) + 6 HCl(aq) → 2 AlCl₃(aq) + 3 H₂O(l) — basic behaviour.
• Al₂O₃(s) + 2 NaOH(aq) + 3 H₂O(l) → 2 NaAl(OH)₄(aq) — acidic behaviour (amphoteric).
• SiO₂(s) + 2 NaOH(l, hot conc) → Na₂SiO₃ + H₂O — acidic; insoluble in water but reacts with hot concentrated alkali.
• P₄O₁₀(s) + 6 H₂O(l) → 4 H₃PO₄(aq) — acidic.
• SO₃(g) + H₂O(l) → H₂SO₄(aq) — strongly acidic.
• Cl₂O₇(l) + H₂O(l) → 2 HClO₄(aq) — strongly acidic.

Explanation

The trend tracks the change from metallic (Na, Mg, Al) to non-metallic (Si, P, S, Cl) elements:

• Na₂O, MgO — formed from metals with low electronegativity. The O²⁻ ions are donated almost completely; the oxide is fully ionic. In water, O²⁻ reacts with H₂O to give OH⁻ (basic).
• Al₂O₃ — the Al³⁺ ion has a high charge density and polarises the O²⁻ significantly, introducing covalent character. The result is amphoteric behaviour: Al₂O₃ reacts with both acids (acting as a base) and alkalis (acting as an acid, forming aluminate ions [Al(OH)₄]⁻).
• SiO₂ — silicon's electronegativity is too high for true ionic bonding; SiO₂ is a giant covalent network of Si–O bonds. It is weakly acidic — insoluble in water but reacts with hot concentrated alkali.
• P₄O₁₀, SO₃, Cl₂O₇ — all molecular covalent oxides. They react with water to form acids of increasing strength (H₃PO₄ weak triprotic acid; H₂SO₄ strong diprotic acid; HClO₄ strong monoprotic acid — one of the strongest acids known).

Worked Example: Predicting Behaviour

Q: Predict whether Al₂O₃ is acidic, basic, or amphoteric, and write equations for any reactions with HCl(aq) and NaOH(aq).

A: Al₂O₃ sits at the boundary between metallic and non-metallic oxides. Aluminium has intermediate electronegativity, and the high charge density of Al³⁺ polarises O²⁻ substantially. Al₂O₃ is therefore amphoteric:

• Al₂O₃(s) + 6 HCl(aq) → 2 AlCl₃(aq) + 3 H₂O(l) (acting as a base)
• Al₂O₃(s) + 2 NaOH(aq) + 3 H₂O(l) → 2 NaAl(OH)₄(aq) (acting as an acid)

The aluminate ion [Al(OH)₄]⁻ is the conjugate base. Amphoteric behaviour is also shown by Al(OH)₃, Zn(OH)₂, BeO, and PbO — a useful periodic-trend pattern.

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Practical Skills Box: Measuring Oxide pH with Universal Indicator

A standard School Science Review-style demonstration links periodicity directly to a measurable laboratory observation. The procedure is straightforward and routinely tested in Section A of Paper 3.

Procedure (qualitative):

1. Place a small spatula-tip of each Period 3 oxide (or a representative subset — typically Na₂O, MgO, Al₂O₃, P₄O₁₀) into separate test tubes containing 10 cm³ of distilled water.
2. Stopper, shake gently, and allow time for dissolution. Al₂O₃ and SiO₂ are essentially insoluble.
3. Add 2 drops of universal indicator to each tube and compare the colour against the pH chart.

Expected observations:

• Na₂O: dissolves rapidly, indicator turns dark purple (pH 13–14 — strongly alkaline).
• MgO: dissolves slowly to give slightly cloudy Mg(OH)₂, indicator turns blue (pH 9–10 — weakly alkaline).
• Al₂O₃: insoluble — no colour change unless reacted with NaOH or HCl in subsequent tests.
• P₄O₁₀: dissolves vigorously (highly exothermic) to give H₃PO₄, indicator turns red (pH 1–2 — strongly acidic).

Safety: Na₂O reacts violently with water and should be added in very small quantities to a large excess of water behind a safety screen. P₄O₁₀ is hygroscopic and corrosive — handle in a fume cupboard. Always wear eye protection.

This practical maps the metallic-to-non-metallic transition onto a single pH ruler and is a memorable way to reinforce the underlying structural argument.

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Summary

Property	Trend across Period 3	Key explanation
Atomic radius	Decreases Na → Cl	Increasing nuclear charge, constant shielding
1st ionisation energy	Generally increases (dips at Al and S)	Increasing Z_eff; sub-shell energy gap (Mg → Al); spin-pair repulsion (P → S)
Electronegativity	Increases Na → Cl	Increasing nuclear charge, decreasing radius
Melting point	Rises across metals (Na, Mg, Al), peaks at Si (giant covalent), falls sharply to P₄, peaks again at S₈, drops to Cl₂ and Ar	Change of structure type dominates over electronic factors
Oxide acidity	Basic → amphoteric → acidic	Bonding character changes ionic → covalent across the period

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Synoptic Links

• atomic-structure L3 (Ionisation Energies) — the successive-IE patterns introduced earlier in the course series provide the direct foundation for understanding the Period 3 first-IE trend. The same arguments about Z_eff, shielding, and sub-shell energy differences are deployed here at the first-IE level.
• L1 of this course (Group 2) — the vertical (down-group) trends in IE, radius, and reactivity for Group 2 are explicitly contrasted with the horizontal (across-period) trends covered here. Together they let you predict the behaviour of any s-block element from its position alone.
• L2 of this course (Group 7) — the halogens (F, Cl, Br, I) show down-group trends opposite in sign to the across-period trends developed here: IE decreases, electronegativity decreases, atomic radius increases. Comparing horizontal and vertical periodicity gives a complete picture of the table.
• L3 of this course (Period 3 Oxides) — the oxide overview given here is expanded in L3 into the full set of physical-state, structural, and acid-base details, with worked equations for the reaction of each oxide with water, acid, and base. Anchors Required Practical 6 (qualitative analysis).
• bonding-and-structure (physical chemistry) — the Period 3 melting-point pattern is the canonical worked example linking macroscopic property (mp) to microscopic structure (metallic, giant covalent, molecular, monatomic). The four-structure framework recurs in every A-Level treatment of structure-property relationships.

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Specimen Exam Question — modelled on the AQA A-Level Chemistry paper format

Question 1. [13 marks total]

(a) Sketch a graph of first ionisation energy (y-axis) against atomic number (x-axis) for the elements Na to Ar. Mark numerical values on the axes and label the two anomalies. [3 marks]

(b) Explain, using electron configurations, why the first ionisation energy of aluminium is lower than that of magnesium. [3 marks]

(c) Silicon has the highest melting point in Period 3. Explain this observation in terms of structure and bonding, and compare with the melting point of phosphorus. [3 marks]

(d) A student dissolves separate samples of Na₂O, P₄O₁₀, and MgO in distilled water, measuring the pH of each resulting solution. Predict the approximate pH of each solution and explain your reasoning in terms of the bonding character of the oxides and the species formed in water. [4 marks]

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Mark Scheme — Specimen Question 1 (modelled, not a real paper)

(a) IE graph across Period 3 [3 marks, AO2]

• 1 mark: general upward trend from Na (≈ 500) through Ar (≈ 1520), with appropriate numerical scale on y-axis.
• 1 mark: explicit dip at Al (lower than Mg) — Al value labelled ≈ 578, Mg ≈ 738.
• 1 mark: explicit dip at S (lower than P) — S value labelled ≈ 1000, P ≈ 1012.

Accept any reasonable scale on the IE axis. Penalise once for a graph showing the dips but with no numbers, or with numbers but no dips.

(b) Mg → Al IE drop [3 marks, AO2]

• 1 mark: electron configurations stated — Mg: [Ne]3s²; Al: [Ne]3s²3p¹ (or equivalent).
• 1 mark: identify that Al's outer electron is in a 3p orbital, which is higher in energy than 3s.
• 1 mark: identify that the 3s² pair provides some additional shielding for the 3p electron, so less energy is needed to remove it from Al than the 3s electron from Mg.

Do not award marks for a vague "Al has a 3p electron" without the energy/shielding argument.

(c) Si melting point [3 marks, AO2]

• 1 mark: identify Si as having a giant covalent (or macromolecular / diamond-like) structure with each Si atom tetrahedrally bonded to four neighbours.
• 1 mark: melting requires breaking many strong Si–Si covalent bonds, hence very high mp.
• 1 mark: phosphorus exists as discrete P₄ molecules with only weak London (dispersion) forces between molecules; melting overcomes only these weak intermolecular forces, giving a much lower mp.

A comparative statement that names the bond/force type broken in each case is required for full marks.

(d) Oxide solutions [4 marks, AO2 + AO3]

• 1 mark: Na₂O is ionic and basic; in water Na₂O + H₂O → 2 NaOH; pH ≈ 13–14.
• 1 mark: MgO is ionic and weakly basic; slightly soluble; gives Mg(OH)₂; pH ≈ 9–10.
• 1 mark: P₄O₁₀ is molecular covalent and acidic; reacts vigorously with water: P₄O₁₀ + 6 H₂O → 4 H₃PO₄; pH ≈ 1–2.
• 1 mark: explicit linkage of acid-base behaviour to bonding character (ionic = basic; molecular covalent = acidic), with reference to the species formed in water (OH⁻ for the metallic oxides; H₃O⁺ from the acidic oxides via H₃PO₄ dissociation).

Accept approximately correct pH ranges. The reasoning mark requires explicit bonding-character logic.

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Grade-Band Model Answers (Question 1)

The three responses below cover the meaningful A-Level range: Grade C (the borderline-pass floor), Grade B (solid mark-scheme coverage), and Grade A* (top-band synthesis). The commentary after each response (editorial, not a real examiner report) names the marks earned and the specific moves that differentiate from adjacent bands.

Grade C response (~280 words)

(a) Graph rises from Na (≈ 500) up to Ar (≈ 1520) with two dips: one at Al (≈ 578) below Mg (≈ 738), and one at S (≈ 1000) below P (≈ 1012). Both dips labelled. Smooth curve through other points.

(b) Mg is [Ne]3s² and Al is [Ne]3s²3p¹. The outer electron in Al is in a 3p orbital, which is higher in energy than the 3s. The 3s electrons in Al also shield the 3p electron slightly. So removing the 3p electron from Al takes less energy than removing the 3s electron from Mg, even though Al has one more proton.

(c) Silicon has a giant covalent structure where every Si atom is bonded to four others by strong covalent bonds in a tetrahedral lattice. To melt it, you have to break many covalent bonds, which takes a lot of energy, so the melting point is very high (1410 °C). Phosphorus exists as P₄ molecules with strong covalent bonds inside the molecule but only weak London forces between molecules. Melting only breaks the weak forces, so the melting point of P₄ is only 44 °C.

(d) Na₂O is ionic and dissolves to give NaOH, so the solution is strongly alkaline (pH ≈ 13). MgO is also ionic but only slightly soluble, giving Mg(OH)₂, so the solution is weakly alkaline (pH ≈ 10). P₄O₁₀ is a molecular covalent oxide and reacts with water to give phosphoric acid H₃PO₄, so the solution is strongly acidic (pH ≈ 2).

Editorial commentary (Grade C): Solid coverage of all four parts with the correct overall structure-bonding framework. To progress to B, the student should quote balanced equations for (d), give explicit shielding arguments in (b), and use the words "London (dispersion) forces" and "intermolecular forces" precisely in (c). The pH predictions are correct but unjustified beyond a single sentence.

Grade B response (~320 words)

(a) Graph rises from Na (496) to Ar (1521) with a clear dip at Al (578, below Mg 738) and at S (1000, below P 1012). Both anomalies labelled. Axes correctly scaled.

(b) Mg: 1s²2s²2p⁶ 3s²; Al: 1s²2s²2p⁶ 3s²3p¹. The outermost electron in Al occupies a 3p orbital, which lies at slightly higher energy than the 3s. In addition, the filled 3s² pair provides a small amount of extra shielding for the 3p electron. Less energy is therefore needed to remove the 3p electron from Al than the 3s electron from Mg, despite Al having one more proton — the sub-shell energy difference outweighs the extra nuclear charge.

(c) Silicon has a giant covalent (macromolecular) structure: each Si atom is covalently bonded to four others in a tetrahedral diamond-cubic lattice, with bond energy 226 kJ mol⁻¹. Melting requires breaking many of these strong covalent bonds, giving a very high mp (1410 °C). Phosphorus exists as discrete P₄ molecules — strong P–P covalent bonds within each tetrahedron but only weak London (dispersion) forces between molecules. Melting overcomes only the weak intermolecular forces (the P–P bonds within P₄ remain intact), so mp = 44 °C — almost two orders of magnitude lower than Si.

(d) Na₂O is ionic; Na₂O + H₂O → 2 NaOH; the OH⁻ ion makes the solution strongly alkaline (pH ≈ 13–14). MgO is ionic but only sparingly soluble; gives Mg(OH)₂; pH ≈ 9–10. P₄O₁₀ is a molecular covalent oxide; P₄O₁₀ + 6 H₂O → 4 H₃PO₄; the H₃PO₄ dissociates as a weak triprotic acid giving pH ≈ 1–2. The pattern (basic → weakly basic → acidic) tracks the change from ionic to covalent bonding across Period 3.

Editorial commentary (Grade B): Now mark-scheme-rigorous: balanced equations in (d), explicit sub-shell + shielding argument in (b), and named force types in (c). To progress to A*, the response could connect the trend to ionic-covalent bonding-character continua, discuss why Al₂O₃ is amphoteric (the "missing oxide" in this question), and reference the polarising power of small high-charge cations.

Grade A* response (~390 words)

(a) Graph: y-axis IE / kJ mol⁻¹ scaled 400–1600, x-axis atomic number Na (11) to Ar (18). Points plotted at Na 496, Mg 738, Al 578, Si 786, P 1012, S 1000, Cl 1251, Ar 1521. Dips at Al and S explicitly annotated with the underlying cause (sub-shell energy and spin-pair repulsion respectively). Connecting line shows the upward trend with the two anomalies as local minima.

(b) Mg: [Ne]3s²; Al: [Ne]3s²3p¹. Two factors lower IE₁(Al) below IE₁(Mg) despite Z increasing by one. First, the 3p orbital lies at slightly higher energy than 3s — penetration of the 3p sub-shell into the inner shells is less effective, so the 3p electron experiences a smaller Z_eff. Second, the filled 3s² pair provides additional shielding for the 3p electron over and above the 1s²2s²2p⁶ core. The net effect is that the 3p electron is more easily removed, lowering IE₁(Al). The same pattern recurs at the s → p boundary in Period 2 (Be → B) and Period 4 (Ca → Ga) — a general periodic feature, not a Period-3 quirk.

(c) Si is a giant covalent solid with the diamond-cubic structure: every Si atom covalently bonded to four neighbours (bond length 235 pm, bond energy 226 kJ mol⁻¹). Melting requires partial disruption of this 3-D bond network, demanding very high temperatures (1410 °C). P exists as discrete P₄ tetrahedra in white phosphorus; strong P–P bonds (200 kJ mol⁻¹) operate within each tetrahedron, but only weak London (dispersion) forces operate between P₄ units. Melting white phosphorus (44 °C) overcomes only these weak intermolecular forces; the strong covalent bonds within each P₄ unit are preserved into the liquid and gas phases. The Si → P drop is therefore not an electronic anomaly but a direct consequence of the change in structure type — from giant covalent network to discrete molecular solid.

(d) Na₂O is fully ionic; the O²⁻ ion is a strong base and reacts with water to give NaOH (pH ≈ 13–14). MgO is also ionic but the Mg²⁺ ion has higher charge density and lower solubility; the saturated solution of Mg(OH)₂ reaches pH ≈ 9–10 only. The intermediate position of Al₂O₃ (amphoteric) reflects the increasing polarising power of the cation: Al³⁺ has sufficient charge density to give the oxide significant covalent character, allowing it to react with both acids and alkalis. P₄O₁₀ is molecular covalent; the electronegativity of P (2.19) is close enough to O (3.44) that the P–O bond is polar covalent rather than ionic, and the oxide reacts vigorously with water: P₄O₁₀ + 6 H₂O → 4 H₃PO₄ (pH ≈ 1–2). The basic-to-acidic trend across Period 3 oxides therefore tracks directly with electronegativity difference Δχ(E–O): large Δχ → ionic → basic; small Δχ → covalent → acidic.

Editorial commentary (Grade A):* Genuinely A*: invokes Z_eff and orbital penetration in (b); quotes bond lengths and energies in (c) and explicitly distinguishes intramolecular from intermolecular forces; in (d) connects the basic-to-acidic trend to electronegativity difference and the polarising power of cations, mentioning Al₂O₃ even though it was not strictly required. The pattern-recognition in (b) (same anomaly recurring at the s → p boundary in every period) demonstrates true synoptic insight.

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Common Errors

• Citing argon in the radius / electronegativity trends. Argon's tabulated radius is a van der Waals radius, not a covalent or metallic radius, and it has no meaningful electronegativity because it does not bond. Restrict comparisons to Na → Cl.
• Explaining the Mg → Al dip without naming both sub-shell-energy and shielding contributions. Both factors are needed for full marks.
• Forgetting "spin-pair repulsion" in the P → S dip. A vague answer about "paired electrons being unstable" earns no marks; the specific phrase "spin-pair repulsion" (or an equivalent precise paraphrase) is required.
• Explaining melting points purely in terms of nuclear charge. The dominant factor is structure type; nuclear-charge arguments apply only within the metallic block (Na, Mg, Al).
• Confusing intramolecular and intermolecular bonds when discussing molecular solids. Melting P₄ or S₈ breaks the London forces between molecules; the strong P–P or S–S covalent bonds within each molecule remain intact.
• Quoting wrong electron configurations. A common slip is writing Al as [Ne]3p¹ (forgetting 3s²). The full notation [Ne]3s²3p¹ is required.

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A-Level Depth Misconceptions

• Shielding is not a binary "on/off" effect. Electrons in the same shell shield each other partially, not fully. Slater's rules (covered in Going Further) quantify this — a 3p electron experiences a Z_eff much lower than Z because the 3s² pair contributes ≈ 0.35 each to the shielding constant.
• The Mg → Al IE drop is small but real and reproducible. It is not a measurement artefact. The same pattern recurs at Be → B, Ca → Ga, and indeed wherever an s-block ends and a p-block begins.
• Periodicity does not stop at melting point. The same Period-3 trend recurs in boiling point, density, electrical conductivity, and thermal conductivity — all driven by the same shift from metallic to covalent to molecular structure.
• The structural change at silicon is the most important fact in Period 3. Without it, the melting-point graph would be a smooth function; with it, you have the characteristic spike at Si and crash at P.
• Oxide acidity is not a single property but a continuum. The shift from basic to acidic is gradual: Al₂O₃ is the canonical amphoteric oxide, but MgO also dissolves slightly in strong acid (so retains a hint of amphoteric character).
• Electronegativity values are conventional, not measured directly. The Pauling scale is derived from bond dissociation energies; the Mulliken scale uses the average of IE and electron affinity; the Allred-Rochon scale uses effective nuclear charge. The Pauling values quoted at A-Level are conventional approximations.

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Going Further

Three undergraduate-adjacent extensions:

• Slater's rules for shielding. A semi-empirical recipe (Slater, 1930) for computing the shielding constant σ from a complete electron configuration, allowing direct calculation of Z_eff for any electron. For a 3p electron in Al: contributions of 0.35 each from the other 3p electrons (in this case none), 0.85 each from the n = 2 shell, and 1.00 each from the n = 1 shell. The resulting Z_eff predicts the observed IE trend semi-quantitatively.
• ZINDO (Zerner's Intermediate Neglect of Differential Overlap) calculations. A modern semi-empirical molecular-orbital method that computes IE, electron affinity, and excitation energies directly from molecular geometry. Period 3 elements are a benchmark test set for the parameterisation. Modern ab initio methods (e.g. DFT with B3LYP functional) give IE values within 1% of experiment.
• Quasi-relativistic effects in heavy elements. As nuclear charge increases down a group, inner-shell electrons approach relativistic speeds, contracting s and p orbitals and destabilising d and f orbitals. The famous example is gold's yellow colour and high IE relative to silver — a direct relativistic effect. Period 3 elements are non-relativistic and the simple Z_eff picture suffices.

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This lesson aligns with AQA A-Level Chemistry §3.2.1 (Periodicity) and provides the foundation for the Group 2 (L1), Group 7 (L2), and Period 3 oxides (L3) treatments that follow. Master the structural framework here — metallic, giant covalent, molecular, monatomic — and the rest of inorganic Period 3 will follow without further memorisation.