Lesson 8: Metallic Bonding and Periodicity

A metal is a lattice of positive ions immersed in a sea of delocalised electrons (Structure 2.3) — electrons that belong to no single atom and move freely through the structure. This one model explains the whole metallic profile: electrical and thermal conductivity (the electrons carry charge and energy), malleability and ductility (layers of ions can slide while the electron sea maintains attraction in every new position — no bonds need breaking in a specific direction), lustre (the mobile electrons absorb and re-emit light), and generally high melting points (the electrostatic attraction between cations and the electron sea is strong). The strength scales with charge density: metals whose atoms release more electrons and have smaller cations bind harder. Periodicity (Structure 3.1) then shows what happens when metallic bonding gives way across Period 3 to giant covalent, then molecular, then monatomic structures — the single most examined pattern in the SL course.

The metallic trend: Na, Mg, Al

Sodium contributes one delocalised electron per atom, magnesium two, aluminium three — and the cations shrink across the period (Na⁺ 102 pm, Mg²⁺ 72 pm, Al³⁺ 54 pm). Both effects strengthen the bonding, and the melting points respond: Na 98 °C, Mg 650 °C, Al 660 °C, with electrical conductivity rising in step (Al is the conductor of choice for overhead cables — light, cheap and highly conductive). Alloys exploit controlled disruption: introducing different-sized atoms (carbon in iron to make steel, zinc in copper to make brass) distorts the layers so they cannot slide easily — harder and stronger, at some cost to conductivity. Pure gold is too soft for jewellery; 18-carat gold is alloyed with copper and silver for exactly this reason.

Period 3: the grand tour of structures

Silicon peaks in melting point because melting a giant covalent network means breaking strong covalent bonds; after silicon the elements form discrete molecules held by feeble dispersion forces, so melting points collapse — sulfur's S₈ rings are bigger than P₄ tetrahedra and Cl₂, so S > P > Cl; argon, with single atoms, is lowest. Electrical conductivity mirrors the bonding story: metals conduct, silicon is a semiconductor (delocalised electrons become available on warming — the foundation of all electronics), and the rest are insulators. The oxides complete the pattern with an acid–base sweep: Na₂O and MgO are basic (Na₂O + H₂O → 2NaOH; suspensions of MgO sit near pH 10), Al₂O₃ is amphoteric — reacting with both acids (Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O) and alkalis (Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄]) — while SiO₂, P₄O₁₀ and SO₂/SO₃ are acidic, the later two forming acids with water (P₄O₁₀ + 6H₂O → 4H₃PO₄; SO₃ + H₂O → H₂SO₄, the chemistry of acid rain).

Worked example 1: explaining the Mg/Na contrast with the model

Explain why magnesium has a higher melting point and better electrical conductivity than sodium. Both are metals: a lattice of cations in delocalised electrons. Magnesium atoms release two electrons per atom against sodium's one, and Mg²⁺ is smaller than Na⁺, so the electrostatic attraction between the cation lattice and the electron sea is stronger in magnesium [model applied]. Melting must weaken this attraction throughout the solid, so magnesium needs more thermal energy — higher melting point. Conductivity follows from carrier density: magnesium's electron sea has twice as many mobile charge carriers per atom, so more charge flows per second under the same potential difference. Every metallic "explain" answer runs on these two rails — number/size of charges for strength, number of mobile electrons for conductivity.

Worked example 2: the oxides practical — pH as the readout

Add a spatula of each Period 3 oxide to distilled water (with warming where needed), shake, and test with universal indicator: Na₂O gives pH ≈ 13–14 (violent reaction — teacher demonstration scale only), MgO ≈ 10, Al₂O₃ and SiO₂ do not dissolve and the water stays ≈ 7 (their character shows only on reaction with strong acid/alkali — Al₂O₃ dissolves in both, the amphoteric test), P₄O₁₀ fizzes to a pH ≈ 1–2 solution, SO₂ (from a gas generator or burning sulfur) gives pH ≈ 3–4. The practical reasoning matters: "insoluble, pH 7" does not mean "neutral oxide" — SiO₂ is acidic but kinetically silent in cold water, as generations of students standing on sandy beaches can confirm. The evaluation point is clean: a single pH test cannot classify an insoluble oxide; the acid/alkali solubility tests complete the classification. Safety thinking likewise: Na₂O and P₄O₁₀ are corrosive and violently hydrolysed, SO₂ is a respiratory hazard — small scale, fume cupboard, goggles, and the reasons stated, not just listed.

Nature of Science — and the real world

The periodic table is chemistry's great classification model — Mendeleev arranged elements by properties and predicted undiscovered ones (gallium, germanium) from the gaps, predictions confirmed within his lifetime. Periodicity works because the underlying variable — electron configuration — is periodic. In the real world this topic is infrastructure: aluminium alloys build aircraft, silicon's semiconductivity builds computers, sulfur oxides from combustion make the acid rain that environmental chemistry fights, and amphoteric aluminium compounds treat both excess stomach acid and acidic lakes.

Common mistakes

• Explaining metallic malleability by "weak bonds" — the bonds are strong; the point is that they are non-directional, so layers slide without breaking them.
• Stating that silicon has a high melting point because of "strong intermolecular forces" — Si is giant covalent; melting breaks covalent bonds.
• Writing Al₂O₃ as simply basic or acidic: it is the SL course's flagship amphoteric oxide, and it must react with both in your answer.
• Forgetting that Ar exists as single atoms: "molecular argon" and "Ar₂" both appear in scripts and both are wrong.

IB Exam Experience

Question 1. Which Period 3 oxide is amphoteric?

• (A) Na₂O
• (B) MgO
• (C) Al₂O₃
• (D) SiO₂

Question 2. Which Period 3 element has the highest melting point?

• (A) Al
• (B) Si
• (C) S
• (D) P

Question 3. Why is steel harder than pure iron?

• (A) Carbon atoms strengthen the iron–iron covalent bonds
• (B) The different-sized carbon atoms distort the lattice layers so they cannot slide over each other easily
• (C) Carbon increases the number of delocalised electrons
• (D) Steel contains ionic bonds between iron and carbon

Answer key. 1. (C) — aluminium oxide reacts with both acids and alkalis. 2. (B) — silicon's giant covalent network requires breaking strong covalent bonds throughout the structure. 3. (B) — alloying atoms of different size pin the layers; there are no covalent or ionic bonds involved.

Structured question (Paper 2 style). Consider the elements sodium to sulfur in Period 3. (a) State the trend in melting point from sodium to aluminium, and explain it in terms of bonding. [3] (b) Explain why silicon has a much higher melting point than phosphorus. [2] (c) Write the equation for the reaction of sodium oxide with water and state the pH of the resulting solution. [2] Marking notes. (a) Melting point increases [1]; the number of delocalised electrons per atom increases (1, 2, 3) and cation size decreases [1]; so the attraction between cations and the electron sea strengthens, requiring more energy to break down the metallic lattice [1]. (b) Silicon is giant covalent — melting breaks strong covalent bonds throughout the network [1]; phosphorus exists as P₄ molecules with only weak dispersion forces between them [1]. (c) Na₂O + H₂O → 2NaOH [1]; pH ≈ 13–14 (accept 12–14) [1].

IA idea: Investigate how the electrical conductivity of a series of metal wires (or alloy samples such as nichrome and brass) varies with composition and temperature, linking the data to the delocalised-electron model.

The groups at the edges: alkali metals and halogens

Periodicity is sharpest inside the groups. The alkali metals share the ns¹ configuration and a personality: soft, low-density metals that react with water to give hydrogen and an alkaline solution — 2Na + 2H₂O → 2NaOH + H₂ — with vigour rising down the group as the outer electron cheapens (Lesson 4's ionisation ledger). The halogens (ns²np⁵) are one electron short and know it: coloured, diatomic, toxic oxidising agents whose reactivity *falls* down the group (F₂ > Cl₂ > Br₂ > I₂) because the incoming electron is accepted into a shell ever further from the nucleus and ever better shielded — the mirror image of the metals' trend, and a standard "explain the opposite trends" pairing. Displacement reactions rank them experimentally: chlorine water added to potassium bromide solution turns orange (Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂ — chlorine is the stronger oxidising agent), while bromine added to potassium chloride does nothing. Add an organic solvent and the halogen's colour partitions into it (Br₂ orange, I₂ violet), sharpening the observation.

The two groups meet in the period's ionic archetypes (NaCl and family, Lesson 5) and in the chlor-alkali industry, which electrolyses brine to chlorine, hydrogen and sodium hydroxide — three commodity chemicals from salt and electricity, and an industrial anchor for Lesson 15's electrolysis. One more trend completes the examinable set: the thermal stability of Group 2 carbonates increases down the group (MgCO₃ decomposes far more readily than BaCO₃) because the small, highly charged-density Mg²⁺ polarises and weakens the carbonate ion most — charge density again, the same variable that ordered the lattice enthalpies in Lesson 5's top-up and the metallic melting points at this lesson's start. Periodicity, fully understood, is a small set of variables — radius, charge, charge density, shielding — recombined; the table is the summary, not the substance.

A first look over the d-block fence

Between aluminium and silicon's group sits the d-block: the transition metals, whose chemistry the HL course revisits and SL students should recognise on sight. Their signatures all trace to one fact — the (n−1)d and ns subshells are close in energy, so d electrons participate variably. Hence variable oxidation states (iron is Fe²⁺ and Fe³⁺; manganese runs from +2 to +7 across its compounds); coloured compounds (partially filled d subshells absorb visible light as electrons move between split d orbitals — Cu²⁺ blue, MnO₄⁻ purple, Cr₂O₇²⁻ orange, each colour a Lesson 3 transition in a new costume); catalysis (variable oxidation states and surfaces that adsorb reactants: iron in Haber, V₂O₅ in the Contact process, nickel in hydrogenation); and complex ions, where ligands donate lone pairs into the metal's empty orbitals by the dative bonding of Lesson 5 — [Cu(H₂O)₆]²⁺ is the pale-blue species in every copper sulfate bottle. The IB's working definition excludes zinc and scandium from "transition metal" status (their common ions have full or empty d subshells, d¹⁰ and d⁰) — a definition with consequences, since Zn²⁺ solutions are colourless for exactly that reason. Keep this map in mind: at HL, Lessons 21 and 25 will walk this territory properly.