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This synoptic lesson ties together the four bonding regimes developed across lessons 0–5: ionic, simple molecular covalent, metallic, and giant covalent (with hydrogen-bonded molecular substances as a structurally important sub-class). Lessons 0–5 covered each at the atomic scale; here we work one level up — the bulk-property level. You will learn to predict melting and boiling points, electrical conductivity in solid/molten/aqueous states, solubility in polar and non-polar solvents, and mechanical properties (hardness, brittleness, malleability) directly from the bonding type, and to run the inference in reverse: given properties, deduce the bonding. This bidirectional structure-property reasoning is among the most heavily examined synoptic skills in AQA Paper 1 and underpins essentially every later topic — energetics, periodicity, redox, transition metals and organic.
Spec mapping (AQA 7405): This lesson sits at the centre of §3.1.3 (bonding) and explicitly draws together the four bonding-type lessons (lessons 0–5) of this course. It feeds into §3.1.4 (energetics — lattice enthalpy quantifies ionic-bond strength; enthalpy of vaporisation reflects intermolecular forces) and §3.2.4 (Period 3 chlorides and oxides provide the canonical worked-example dataset for bonding-type trends). It also connects to §3.3.1 (IUPAC nomenclature — once you can name a compound systematically, you can predict its properties from its formula and inferred bonding). Refer to the official AQA specification document for exact wording.
Assessment objectives: Recall of the property-type mapping for the four bonding categories is AO1. Applying the mapping in either direction — given the bonding type, predict the properties; given the properties, deduce the bonding type — is AO2 and appears on every Paper 1. Rationalising borderline cases (polar covalent vs ionic with partial covalent character, metalloid behaviour, SiO₂ as giant covalent vs SO₂ as simple molecular) is AO3 and discriminates the top mark band.
Before predicting properties, fix the canonical examples and the structural picture for each bonding regime. Lessons 0–5 developed each in detail; the table below is a compressed recap and the anchor for the property-prediction work that follows.
| Bonding type | Canonical example | Particles in the lattice | Forces holding particles together |
|---|---|---|---|
| Ionic | NaCl | Cations and anions in a 3-D giant lattice | Strong electrostatic attraction between full opposite charges |
| Simple molecular covalent | HCl, I₂, CO₂ | Discrete molecules in a 3-D molecular lattice | Strong covalent bonds within each molecule; weak intermolecular forces (London dispersion, permanent dipole, hydrogen bonding) between molecules |
| Metallic | Cu, Na, Al | Cations in a 3-D giant lattice, immersed in a sea of delocalised valence electrons | Strong electrostatic attraction between cations and the delocalised electron sea |
| Giant covalent | Diamond, graphite, SiO₂, Si | Atoms covalently bonded in a continuous 3-D network (or 2-D sheets for graphite) | Strong covalent bonds throughout — there is no separation between the particle and the lattice |
The critical structural distinction for physical properties is between giant structures (ionic, metallic, giant covalent — where the strong bonds form a continuous lattice) and simple molecular structures (where the strong covalent bonds are confined to the inside of each molecule, and only weak intermolecular forces hold the molecules together in the bulk solid or liquid).
When a substance melts, the particles gain enough kinetic energy to overcome the forces that hold them in their lattice positions. When it boils, they overcome the forces holding them in the liquid altogether and escape into the gas phase. The question to ask for any substance is therefore: what forces have to be broken on melting and boiling?
For ionic substances, melting breaks ionic bonds — the full electrostatic attractions between cations and anions throughout the lattice. These are very strong (typically 600–4000 kJ mol⁻¹ of lattice enthalpy), so melting points are high. NaCl melts at 801 °C; MgO at 2852 °C. The trend across ionic compounds tracks the magnitude of the charges and the inverse of the inter-ion distance (Coulomb's law) — a quantitative connection developed in the §3.1.4 energetics chapter.
For metallic substances, melting weakens the cation–electron-sea attraction enough that the cations can flow past one another, but it does not fully separate them; full separation occurs only on boiling. Melting points are moderate-to-high (Hg −39 °C, Na 98 °C, Cu 1085 °C, W 3422 °C) and increase strongly with the charge on the cation and with the number of delocalised electrons per atom (Group 1 cations have charge +1 and contribute one electron; Group 2 contribute two; transition metals contribute several).
For giant covalent substances, melting requires breaking the covalent bonds themselves — and these are the strongest of all the bond types here (around 350 kJ mol⁻¹ per C–C bond in diamond). Melting points are accordingly extreme: diamond 3550 °C, SiO₂ 1610 °C, graphite sublimes around 3650 °C.
For simple molecular substances, melting does not break the covalent bonds inside each molecule — those are far too strong to break thermally below several hundred degrees. Instead, melting and boiling break only the intermolecular forces between molecules. These are weak (typically 1–30 kJ mol⁻¹), so melting and boiling points are low. HCl boils at −85 °C; CO₂ sublimes at −78 °C; I₂ melts at 114 °C only because its large electron cloud gives it unusually strong London dispersion forces.
The chlorides of the Period 3 elements form the canonical AQA dataset for the transition from ionic through polar covalent to simple molecular bonding. Use this table to anchor the property-prediction ladder:
| Chloride | Bonding | Structure | mp (°C) | bp (°C) |
|---|---|---|---|---|
| NaCl | Ionic | Giant ionic lattice | 801 | 1465 |
| MgCl₂ | Ionic (with some covalent character) | Giant ionic lattice (layered) | 714 | 1412 |
| AlCl₃ | Polar covalent | Layered molecular lattice; sublimes at 192 °C; dimerises in vapour as Al₂Cl₆ | 192 (sub.) | — |
| SiCl₄ | Simple molecular covalent | Tetrahedral molecule | −68 | 57 |
| PCl₃ | Simple molecular covalent | Trigonal pyramidal molecule | −94 | 76 |
| SCl₂ | Simple molecular covalent | Bent molecule | −121 | 59 |
| Cl₂ | Simple molecular covalent | Diatomic molecule | −101 | −34 |
The melting points fall sharply between MgCl₂ (ionic) and AlCl₃ (polar covalent), and again between AlCl₃ and SiCl₄ (simple molecular). Note that there is no giant covalent chloride at Period 3 — silicon chloride is SiCl₄ (a discrete molecule), not a network. Silicon's giant covalent regime appears in its oxide SiO₂ (mp 1610 °C), not in its chloride. This is the cleanest illustration in the AQA syllabus of how the same element (Si) can sit in completely different bonding regimes in different compounds, and how the bulk physical properties follow.
Electrical conduction requires mobile charge carriers — either electrons or ions. The conductivity of a substance therefore depends not only on the bonding type but also on the physical state, because state changes can mobilise or immobilise the carriers. The following 4 × 3 table is the central object to memorise.
| Bonding type | Solid | Molten | Aqueous |
|---|---|---|---|
| Ionic | Insulator | Conductor (mobile ions) | Conductor (dissolved ions) — if soluble |
| Metallic | Conductor (delocalised electrons) | Conductor (delocalised electrons) | Not applicable (insoluble); some react with water |
| Simple molecular covalent | Insulator | Insulator | Insulator — unless the molecule autoionises in water (HCl, NH₃) |
| Giant covalent | Insulator | Insulator | Insoluble — n/a |
Key teaching points:
Common Misconception: "Ionic compounds conduct because they contain ions." Incomplete — the ions must also be mobile. Solid NaCl contains ions but they are not free to move; it is an insulator. Molten NaCl contains the same ions in mobile form and conducts. Always specify the state when discussing ionic conductivity.
The qualitative rule is "like dissolves like" — polar solvents dissolve polar (and ionic) solutes; non-polar solvents dissolve non-polar solutes. Dissolution requires solute–solute and solvent–solvent attractions to be replaced by solute–solvent attractions of comparable energy; when polarities match, the balance is favourable.
Key Point: "Insoluble" at A-Level means "not significantly soluble in water at room temperature". Strictly, every solid has a finite (often vanishingly small) solubility; the qualitative rule is enough for most predictions.
Mechanical properties trace back to what happens when atomic layers are pushed past one another by an applied force.
In Paper 1 you will frequently be given a property profile and asked to identify the bonding type. The workflow is: (1) look at the melting point (giant lattice or simple molecular?); (2) look at the conductivity (delocalised electrons, mobile ions, or neither?); (3) look at the solubility and mechanical properties to confirm. The following four worked examples cover the four main bonding regimes.
Example 1. White crystalline solid; mp 808 °C; insulator as solid; conductor when molten; soluble in water giving a conducting solution; brittle with clean cleavage planes.
Reasoning: High mp rules out simple molecular. Insulator solid → conductor melt → mobile carriers appear only when the lattice is broken → ionic. Aqueous conductivity (solvated ions) and brittleness (layer slip → like-charge repulsion → shatter) confirm. Consistent with NaCl (mp 801 °C).
Example 2. Silvery solid; mp 660 °C; conductor as solid and molten; insoluble in water and organic solvents; malleable.
Reasoning: Conductor in both solid and molten state → mobile electrons in all states → metallic. Malleability confirms (cations slip without like-charge repulsion as the electron sea redistributes). Matches aluminium (mp 660 °C).
Example 3. Colourless, glassy solid; mp 1610 °C; insulator solid and molten; insoluble; extremely hard; main component of sand.
Reasoning: Very high mp → giant lattice. Insulator in all states → not ionic, not metallic. Extreme hardness and insolubility → valence electrons locked into covalent bonds throughout a 3-D network → giant covalent. Identification: SiO₂.
Example 4. Colourless liquid; mp −95 °C, bp 69 °C; insulator in all states; insoluble in water but miscible with hydrocarbons.
Reasoning: Very low mp and bp → only weak IM forces being broken → simple molecular covalent. Insulator throughout → no mobile carriers. "Like dissolves like" with non-polar solvents → non-polar molecule. Matches hexane, C₆H₁₄.
Exam Tip: The first thing to ask is always: "does the substance have a giant or a simple molecular lattice?" The melting point tells you immediately — under about 300 °C strongly suggests simple molecular; over about 500 °C strongly suggests giant. Then use conductivity and solubility to discriminate between the three giant types (ionic / metallic / giant covalent).
Real substances do not always sit cleanly inside one of the four idealised categories. The most important borderline cases at A-Level are:
The four bonding categories are a useful classification, but the underlying physics is a continuum — Fajans's rules and the van Arkel–Ketelaar triangle (Going Further) place each compound on it.
Hydrogen bonding is an unusually strong intermolecular force (typically 10–40 kJ mol⁻¹, vs 1–10 kJ mol⁻¹ for ordinary London dispersion). It occurs when an H atom bonded to N, O or F interacts with a lone pair on another such atom. The consequences for physical properties:
Hydrogen-bonded substances remain simple molecular — the IM forces are stronger than usual but still far weaker than the bonds in a giant lattice.
| Bonding type | Melting point | Solid conductivity | Molten conductivity | Aqueous conductivity | Solubility in water | Hardness |
|---|---|---|---|---|---|---|
| Ionic | High (600–3000 °C) | Insulator | Conductor | Conductor (if soluble) | Often soluble | Hard but brittle |
| Polar molecular | Low–moderate (−100 to 200 °C) | Insulator | Insulator | Conductor only if ionises (e.g. HCl) | Variable; often soluble | Soft |
| Non-polar molecular | Very low (−200 to 100 °C) | Insulator | Insulator | Insulator (insoluble) | Insoluble | Soft |
| Metallic | Moderate–very high (−39 to 3400 °C) | Conductor | Conductor | n/a (insoluble) | Insoluble (or reactive) | Malleable, ductile |
| Giant covalent | Very high (1000–4000 °C) | Insulator (graphite exception) | Insulator | n/a (insoluble) | Insoluble | Very hard (diamond) or soft along layers (graphite) |
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