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Metallic bonding describes the cohesion of a metallic solid as a regular lattice of positive cations immersed in a delocalised "sea" of valence electrons. This single picture — cations held in place by a shared, mobile electron cloud — accounts for the entire familiar suite of metallic properties: high melting points, electrical and thermal conductivity even in the solid state, malleability, ductility and lustrous appearance. The same model rationalises why alloys are typically harder than the pure metals they contain. In the second half of this lesson we contrast metallic bonding with the giant covalent (macromolecular) structures of diamond, graphite, graphene, silicon and silicon dioxide. These materials are held together by directional, localised covalent bonds, yet their physical properties span the entire spectrum from hard, transparent insulator (diamond) through soft conducting solid (graphite) to room-temperature semiconductor (silicon). Mastering the contrasts between these structure types is the central deliverable of AQA §3.1.3.
Spec mapping (AQA 7405): This lesson maps to §3.1.3 (bonding) — specifically the metallic-bonding model, properties of metals, alloys, and giant covalent structures (diamond, graphite, graphene, SiO₂). It builds on lesson 0 (ionic bonding and giant ionic lattices) and lesson 1 (covalent and dative bonding), is synthesised in lesson 6 (physical properties and bonding types — the explicit comparison of all bonding regimes), and feeds forward to §3.2.5 (transition metals, where partially-filled d-orbitals contribute additional electrons to the metallic-bonding picture) and §3.1.4 (energetics, where the standard enthalpy of atomisation ΔH°_at gives a direct quantitative measure of metallic-bond strength). Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: AO1 recall items include the definition of a metallic bond, the listing of metallic properties (electrical and thermal conduction, malleability, ductility, high melting point, lustre) and the structural facts for diamond, graphite, graphene and SiO₂. AO2 application requires predicting and comparing metallic-bond strength from cation charge, cation radius and the number of delocalised valence electrons (the canonical Na/Mg/Al sequence), explaining alloy hardness in terms of disrupted slip planes, and explaining why graphite conducts but diamond does not. AO3 demands rationalising property contrasts across the four giant covalent allotropes, predicting semiconductor vs insulator behaviour from a qualitative band-gap argument, and constructing extended-prose comparisons (often 6-marker style) that integrate evidence from melting points, conductivity and mechanical behaviour.
Key Definition: A metallic bond is the strong electrostatic attraction between a regular lattice of positive metal cations and the delocalised valence electrons that surround them.
In a metallic solid, each atom loses its outer (valence) electrons. These electrons no longer "belong" to any particular atom — they are free to move throughout the entire macroscopic crystal. The atoms that remain are positively charged cations, arranged in a regular three-dimensional close-packed lattice (face-centred cubic, body-centred cubic or hexagonal close-packed, depending on the metal). The bond is the net electrostatic attraction between every cation and every nearby delocalised electron. It is non-directional — unlike a covalent bond, the metallic bond points everywhere at once because the electron cloud is isotropic.
+ e⁻ + e⁻ +
e⁻ + e⁻ + e⁻
+ e⁻ + e⁻ +
e⁻ + e⁻ + e⁻
+ e⁻ + e⁻ +
Historically this picture is associated with the early-twentieth-century work of Drude (who proposed a free-electron gas model of metals) and its quantum-mechanical refinement by Sommerfeld. At A-Level you only need the simple "sea of electrons" cartoon, but it is worth knowing that the rigorous treatment replaces the classical electron gas with a band structure — valence and conduction bands that overlap in metals (we return to this in Going Further). For all standard exam questions, the cation-sea language is sufficient and correct.
Three factors determine the strength of the metallic bond, and therefore the melting point, boiling point and standard enthalpy of atomisation of the metal:
| Factor | Effect on metallic-bond strength | Example |
|---|---|---|
| Cation charge | Higher charge → stronger electrostatic attraction to the electron sea | Mg²⁺ > Na⁺; Al³⁺ > Mg²⁺ |
| Cation radius | Smaller cation → electron sea is closer → stronger attraction | Na > K > Rb (down Group 1) |
| Number of delocalised electrons per atom | More delocalised electrons → denser electron cloud → stronger bond | Al (3e⁻) > Mg (2e⁻) > Na (1e⁻) |
These factors are not independent — across Period 3 from Na → Mg → Al, the nuclear charge increases, the atomic radius contracts and the number of delocalised electrons rises in step. All three effects pull in the same direction, producing a steep rise in melting point.
| Metal | Outer electron config. | Delocalised e⁻ per atom | Cation | Cation radius (pm) | Melting point (°C) |
|---|---|---|---|---|---|
| Na | 3s¹ | 1 | Na⁺ | 102 | 98 |
| Mg | 3s² | 2 | Mg²⁺ | 72 | 650 |
| Al | 3s² 3p¹ | 3 | Al³⁺ | 54 | 660 |
The jump from Na to Mg is dramatic (98 → 650 °C, a factor of ~6.6 in absolute terms): doubling the cation charge from +1 to +2, halving the cation radius and doubling the number of delocalised electrons all reinforce each other. The further rise to Al is modest in melting-point terms, but the boiling point and the standard enthalpy of atomisation (Na 107, Mg 148, Al 326 kJ mol⁻¹) confirm the continuing trend in bond strength.
Group 1 (Na → K → Rb → Cs) shows monotonically falling melting points (98 → 64 → 39 → 28 °C) as the cation grows and the electron sea is pulled further from the nucleus — exactly as the radius factor predicts. Transition metals such as Fe (1538 °C), W (3422 °C) and Os (3033 °C) have very high melting points because they contribute partially-filled d-electrons to the delocalised sea, generating additional bonding density. Mercury, conversely, melts at −39 °C — a relativistic-contraction effect on the 6s electrons that we do not pursue here, but it is a memorable counter-example to the simple periodic trend.
Key Point: When asked to compare metallic-bond strength, name all three factors (charge, radius, electron count) and identify which dominate in the case at hand. Single-factor answers rarely score full AO2 marks.
Every familiar metallic property follows directly from the cation-sea picture.
| Property | Explanation rooted in the cation-sea model |
|---|---|
| High melting and boiling points | Many strong, non-directional electrostatic attractions between cations and delocalised electrons; large amounts of energy required to disrupt the lattice. |
| Good electrical conductivity (solid and liquid) | Delocalised valence electrons are mobile and respond to an applied potential difference, carrying current through the bulk solid even before melting. |
| Good thermal conductivity | Delocalised electrons rapidly transport kinetic energy from a hot region of the crystal to a cold one; lattice vibrations (phonons) contribute but the electron contribution dominates in pure metals. |
| Malleability (hammered into sheets) | Layers of cations can slip past one another without rupturing bonds, because the non-directional electron sea reorganises to surround the cations in their new positions. |
| Ductility (drawn into wires) | Same mechanism — slip of close-packed layers under tension is accommodated by the mobile electron cloud. |
| Lustrous (shiny) appearance | Delocalised electrons absorb visible photons across a continuous range of frequencies and immediately re-emit them; the surface acts as a near-perfect reflector across the visible spectrum. |
| High density | Close-packing of cations leaves very little void space; metals such as Os (22.6 g cm⁻³) and Ir (22.5 g cm⁻³) are among the densest known elements. |
| Metal | Melting point (°C) | Comment |
|---|---|---|
| Hg | −39 | Liquid at room temperature (relativistic effect on 6s electrons) |
| Cs | 28 | Lowest melting Group 1 metal |
| Na | 98 | Soft, low charge, large radius |
| Mg | 650 | +2 cation, two electrons per atom |
| Al | 660 | +3 cation, three electrons per atom |
| Cu | 1085 | d-electrons contribute |
| Fe | 1538 | Transition metal |
| W | 3422 | Highest-melting metallic element |
Key Contrast — Metals vs Ionic Solids: Both have giant lattices and high melting points, but metals are malleable and conduct in the solid state; ionic solids are brittle and conduct only when molten or in solution. The difference traces directly to the mobile electron sea in metals versus the rigid, directional Coulombic lattice in ionic solids.
Key Definition: An alloy is a solid solution of two or more elements, at least one of which is a metal, with metallic properties. It is not a compound — there is no fixed stoichiometry, and the components are not joined by ionic or covalent bonds.
In a pure metal, every cation is identical. Layers of close-packed cations slide easily over one another when a shear stress is applied — this is the microscopic origin of malleability and ductility. In an alloy, atoms of a different element (a different size from the host metal's cations) sit within the lattice and disrupt the regular slip planes. The misfit atoms pin the dislocations that would otherwise mediate slip, raising the stress required to deform the material.
Pure metal (slip easy): Alloy (slip pinned):
○ ○ ○ ○ ○ ○ ○ ○ ○ ○
○ ○ ○ ○ ○ ○ ● ○ ○ ○ ● = different-sized atom
○ ○ ○ ○ ○ ○ ○ ○ ● ○
○ ○ ○ ○ ○ ○ ○ ○ ○ ○
The hardening effect is generally largest when the size mismatch is largest, up to a limit beyond which the misfit atom no longer dissolves in the lattice and a separate phase precipitates instead. The size-mismatch principle is why steel (carbon in iron) is much harder than pure iron, and why brass (zinc in copper) is harder than pure copper.
| Alloy | Composition | Key property | Use |
|---|---|---|---|
| Brass | Cu + Zn (~30%) | Hard, corrosion-resistant, gold-coloured | Decorative, musical instruments, plumbing fittings |
| Bronze | Cu + Sn (~12%) | Hard, corrosion-resistant | Historical weaponry, statues, bearings |
| Steel | Fe + C (0.05–2%) | Hard, strong, ductile | Construction, vehicles, tools |
| Stainless steel | Fe + Cr (~18%) + Ni (~8%) | Hard, corrosion-resistant (Cr₂O₃ passivation layer) | Kitchenware, surgical instruments, cladding |
| Solder | Sn + Pb (or Sn + Cu/Ag) | Low melting point | Joining electrical components |
| Duralumin | Al + Cu (~4%) + Mg + Mn | High strength-to-weight | Aircraft frames |
Despite the disruption of slip planes, alloys retain the qualitative metallic-bonding picture — the electron sea persists across both host and guest atoms — so they still conduct electricity, conduct heat and have metallic lustre.
Giant covalent structures are continuous three-dimensional or two-dimensional networks of atoms joined by directional, localised covalent bonds. Unlike simple molecular substances (which have discrete molecules held together by weak intermolecular forces), giant covalent solids have no molecules at all — the entire crystal is one macroscopic supermolecule. Melting therefore requires breaking many strong covalent bonds, giving very high melting and sublimation points across the class.
Structure: Each carbon atom is sp³ hybridised and bonded to four neighbouring carbons by single C—C σ bonds in a perfect tetrahedral arrangement. The C—C bond length is 154 pm; the bond angle is 109.5°. The lattice extends in three dimensions; every electron is localised in a bonding orbital.
| Property | Value / explanation |
|---|---|
| Sublimation point | ~3550 °C (diamond sublimes rather than melts at standard pressure) |
| Hardness (Mohs) | 10 — hardest naturally occurring substance |
| Electrical conductivity | Insulator — all valence electrons are in σ bonds; no free charge carriers |
| Thermal conductivity | Very high — phonon (lattice-vibration) transport is highly efficient in a rigid 3D network |
| Optical | Transparent in the visible region; very high refractive index (n ≈ 2.42) |
Structure: Each carbon is sp² hybridised and bonded to three neighbouring carbons by σ bonds in a flat hexagonal arrangement. The bond angle within a layer is 120°; the in-plane C—C bond length is 142 pm (shorter than in diamond — bond order is intermediate between 1 and 2 because of the delocalised π system). The fourth valence electron on each carbon occupies a p-orbital perpendicular to the layer, and these p-orbitals overlap sideways to form a delocalised π system spread across the entire sheet. Sheets are stacked ~335 pm apart and held together by weak London dispersion forces.
| Property | Value / explanation |
|---|---|
| Sublimation point | ~3700 °C — within-sheet σ bonds remain very strong |
| Hardness | Soft and slippery — sheets slip past one another under shear |
| Electrical conductivity | Conductor within the plane (delocalised π electrons mobile); much poorer perpendicular to the planes |
| Lubricant behaviour | Sheets slide past one another due to weak interlayer London forces |
| Common use | Pencils (sheets transfer to paper); electrodes; high-temperature crucibles |
Exam Tip: Always state that graphite conducts because each carbon contributes one electron to a delocalised π system spread across the sheet — not because "the structure is layered". The layering is responsible for the softness, not the conductivity.
Structure: A single isolated sheet of graphite — a one-atom-thick, two-dimensional honeycomb of sp² carbons. The in-plane structure is identical to a single graphite layer, but graphene is studied as a free-standing two-dimensional material with no layers above or below. Graphene was first isolated experimentally in 2004 by Geim and Novoselov, whose work was recognised by the Nobel Prize in Physics in 2010.
| Property | Value / explanation |
|---|---|
| Electrical conductivity | Outstanding — electron mobility ~100× that of silicon at room temperature; the delocalised π system is the largest known |
| Mechanical strength | Tensile strength ~100× that of steel by weight |
| Optical | Transparent — absorbs only ~2.3% of incident visible light per layer |
| Thermal conductivity | Very high in-plane (~5000 W m⁻¹ K⁻¹ — higher than diamond) |
| Areal density | ~0.77 mg per m² — extraordinarily light |
Applications under active research include flexible electronics, high-speed transistors, transparent conducting electrodes, ultrasensitive chemical sensors, supercapacitor electrodes and composite reinforcement. Graphene is also the building block from which carbon nanotubes (rolled-up sheets) and fullerenes (closed cages) can be conceptually constructed.
Structure: Each silicon atom sits at the centre of a tetrahedron of four oxygens; each oxygen bridges two silicons. The empirical formula SiO₂ reflects this 1:2 ratio — there is no discrete SiO₂ molecule. The network extends in three dimensions with strong, polar Si—O σ bonds (Si—O bond enthalpy ~464 kJ mol⁻¹; Si—O bond length ≈ 162 pm).
| Property | Value / explanation |
|---|---|
| Melting point | ~1710 °C — many strong Si—O bonds must break |
| Hardness | Very hard (Mohs 7) — rigid 3D framework |
| Electrical conductivity | Insulator — all valence electrons in localised Si—O bonds |
| Solubility in water | Insoluble — bonds are too strong to be replaced by solvation |
| Optical | Transparent in the visible; quartz is widely used in optics |
Elemental silicon adopts the same diamond-type lattice (each Si bonded to four others tetrahedrally) but with a smaller band gap (~1.1 eV vs ~5.5 eV for diamond). The smaller gap means that at room temperature a small fraction of valence electrons are thermally excited into the conduction band, so silicon is a semiconductor — neither a metallic conductor nor a true insulator. The temperature dependence of silicon's conductivity (rising with increasing T) is opposite to that of a metal (whose conductivity falls with T, because lattice vibrations scatter the conduction electrons more strongly). This contrast is one of the cleanest experimental signatures of band-gap behaviour.
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