HookA pencil and a diamond are the same atoms
In December 1954, at General Electric's laboratory in Schenectady, New York, Tracy Hall squeezed carbon past roughly 100,000 atmospheres at around 1,600 °C in a press of his own design — and made the first verified synthetic diamonds. The starting material was essentially the stuff of pencil lead. Graphite is so soft it smears onto paper, and it conducts electricity; diamond is the hardest natural material known and an electrical insulator. Same element, same atoms, wildly different substance — and GE's reward to Hall for one of the most valuable patents of the century was, famously, a ten-dollar savings bond.
The lesson is the whole of section 3.1.3: properties do not live in atoms; they live in the forces between them. This section hands you the complete toolkit — three types of strong bond (ionic, covalent including dative, metallic), four types of crystal, a geometry engine (electron-pair repulsion theory), a polarity dial (electronegativity) and three intermolecular forces. Every exam question here, from 1-mark definitions to 6-mark property comparisons, is secretly the same question: which particles are present, which forces hold them, and how much energy does it take to pull them apart?
ModelIonic bonding: charged spheres and the lattice they build
An ionic bond is the electrostatic attraction between oppositely charged ions — that exact phrase, because AQA pays for the word electrostatic. The ions form by electron transfer from metal to non-metal, and because the attraction acts equally in every direction, the result is not a molecule but a giant ionic lattice: in sodium chloride each Na\(^{+}\) is held by six surrounding Cl\(^{-}\) ions and vice versa, repeating in all directions.
Lattice strength scales with the charges on the ions and falls as the ions get bigger — higher charge density means stronger attraction. Compare melting points: NaCl (1+, 1−) melts at 801 °C, while MgO (2+, 2− and smaller ions) holds out to 2,852 °C, which is why magnesium oxide lines furnaces. You are also expected to write formulas fluently from ions, including the compound ions the spec names: sulfate \(\text{SO}_4^{2-}\), hydroxide \(\text{OH}^{-}\), nitrate \(\text{NO}_3^{-}\), carbonate \(\text{CO}_3^{2-}\) and ammonium \(\text{NH}_4^{+}\). Balance the charges to zero and the formula writes itself: ammonium sulfate must be \((\text{NH}_4)_2\text{SO}_4\).
ModelCovalent and dative covalent: shared pairs, whatever their origin
A covalent bond is a shared pair of electrons between two atoms — each nucleus attracted to the shared pair sitting between them. Single, double and triple bonds are one, two and three shared pairs. Covalent bonding builds discrete molecules (\(\text{H}_2\text{O}\), \(\text{CO}_2\)) and, as the next block shows, entire giant structures (diamond).
A dative covalent (coordinate) bond is a covalent bond in which both electrons of the shared pair come from the same atom — one species supplies a lone pair, the other supplies an empty orbital. Drawn as an arrow from donor to acceptor, it is chemically indistinguishable from any other covalent bond once formed. The set pieces: ammonia donating its lone pair to H\(^{+}\) to make ammonium, \(\text{NH}_3 + \text{H}^{+} \rightarrow \text{NH}_4^{+}\) (all four N–H bonds then identical); water doing the same to make the oxonium ion \(\text{H}_3\text{O}^{+}\); and aluminium chloride dimerising to \(\text{Al}_2\text{Cl}_6\), where a chlorine lone pair on each \(\text{AlCl}_3\) plugs the electron-deficient aluminium of the other.
ModelMetallic bonding: cations in a sea of electrons
A metal is a giant lattice of positive ions sitting in a sea of delocalised electrons — the outer electrons of every atom, surrendered to the whole structure. The bond is the attraction between the cations and that delocalised sea, and it explains the metallic property set in one move: conductivity, because the delocalised electrons are mobile charge carriers; malleability, because layers of ions can slide past each other while the non-directional electron glue keeps holding; and generally high melting points, because the attraction runs through the entire lattice.
Strength is tunable by three factors: the charge on the metal ion (and with it, electrons contributed per atom), and the ionic radius — more charge, more delocalised electrons and smaller ions all mean stronger bonding. Period 3 shows the trend cleanly: sodium (1+) melts at 98 °C, magnesium (2+) at 650 °C, aluminium (3+) at 660 °C and is markedly harder and stronger than either. This is the comparison AQA reaches for when it wants a quick AO2 test of the model.
MechanismFour crystal types: one decision tree for every property question
AQA names four crystal structures, each with a flagship example you should attach permanently: ionic (sodium chloride), metallic (magnesium), macromolecular — also called giant covalent — (diamond and graphite), and simple molecular (iodine, ice). Every melting-point, conductivity or hardness question is the same three-step decision: identify the structure, name the particles and the forces between them, then judge the energy needed to overcome those forces.
Melting separates them brutally. Ionic and metallic lattices need strong electrostatic attractions overcome: NaCl at 801 °C, Mg at 650 °C. Macromolecular structures must break covalent bonds themselves, so diamond survives past 3,500 °C. Simple molecular crystals melt when their weak intermolecular forces let go — the covalent bonds inside each molecule never break — so iodine slips apart at 114 °C and ice at 0 °C. Conductivity is a hunt for mobile charge: metals conduct always (delocalised electrons); ionic compounds only when molten or dissolved, because only then can the ions move; molecular substances never. Graphite is the beautiful exception that proves every rule: each carbon bonds to only three others in layers, the fourth electron per atom delocalises between the layers (so it conducts), and the layers, held only by weak forces, slide — which is precisely why a pencil writes while a diamond drill cuts.
ModelShapes: electron pairs repel, and lone pairs repel hardest
Electron-pair repulsion theory is the geometry engine: pairs of electrons around a central atom repel one another and settle as far apart as possible. Count the pairs and the base shape follows — 2 pairs linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 5 trigonal bipyramidal (120° and 90°), 6 octahedral (90°). A double bond counts as one region of electron density, which is why \(\text{CO}_2\) is linear.
The refinement that earns the extra mark: lone pairs repel more than bonding pairs, because they sit closer to the central atom. Each lone pair squeezes the remaining bond angles by about 2.5°, and the shape is named from the atom positions only. So with four pairs: methane (4 bonding) is tetrahedral at 109.5°; ammonia (3 bonding + 1 lone) is trigonal pyramidal at 107°; water (2 + 2) is bent at 104.5°. The method extends to any ion once you adjust the electron count for the charge — which is exactly what the worked example does.
The transferable routine: count the central atom's outer electrons, add one per single bond, adjust for charge, halve for pairs, then split into bonding and lone. (1) \(\text{AlH}_4^{-}\): aluminium brings 3 electrons, four hydrogens bring 4, the negative charge adds 1, total 8 — four pairs, all bonding. Shape: tetrahedral, 109.5°. (2) \(\text{ClF}_3\): chlorine brings 7, three fluorines bring 3, total 10 — five pairs: three bonding, two lone. Five pairs adopt a trigonal bipyramid, and the two lone pairs take equatorial positions to minimise repulsion, leaving the three atoms in a T-shape with F–Cl–F angles pushed just below 90° (about 87.5°). Write the electron count, the pair split, the shape name and the angle — that sequence is the mark scheme, in order.
ModelElectronegativity: the tug-of-war that makes bonds polar
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond, measured on the Pauling scale. Fluorine tops the table at 4.0, with oxygen 3.5, nitrogen and chlorine about 3.0, and hydrogen 2.1. It increases across a period (nuclear charge rises, radius shrinks) and decreases down a group (shielding and distance win). When two bonded atoms differ in electronegativity, the shared pair sits closer to the more electronegative atom, leaving partial charges \(\delta+\) and \(\delta-\): a polar bond. Push the difference far enough and the bond is better described as ionic — covalent and ionic bonding are ends of one continuum, not separate boxes.
A polar molecule needs more than polar bonds: the bond dipoles must fail to cancel. Symmetry is the executioner — carbon dioxide's two C=O dipoles pull in exactly opposite directions and cancel (linear), and tetrahedral \(\text{CCl}_4\) cancels four ways, so both molecules are non-polar despite strongly polar bonds. Bent water, with its dipoles at 104.5°, does not cancel — and that permanent dipole underwrites most of the next block.
MechanismThree intermolecular forces — and why ice floats
Between molecules act three forces, in rising order of strength. Induced dipole–dipole forces (van der Waals dispersion forces) exist between all molecules: electrons move randomly, an instantaneous dipole flickers into being, and it induces dipoles in neighbours. They grow with the number of electrons and the area of contact — which is why the alkanes' boiling points climb with chain length, why branched isomers boil lower than straight ones, and why a gecko can hang its whole body weight from a pane of glass on nothing but dispersion forces between the glass and half a million microscopic hairs per foot. Permanent dipole–dipole forces add an extra attraction between polar molecules such as propanone. Hydrogen bonding — the strongest of the three, though still roughly a tenth the strength of a covalent bond — needs a specific arrangement: hydrogen bonded directly to nitrogen, oxygen or fluorine, attracted to a lone pair on the N, O or F of a neighbouring molecule.
The evidence is on a graph AQA loves: the boiling points of \(\text{NH}_3\), \(\text{H}_2\text{O}\) and HF tower over the trend of their groups — water boils at 100 °C while hydrogen sulfide, twice the mass, boils at −60 °C. Hydrogen bonding also explains ice. Each water molecule hydrogen-bonds to four others in an open hexagonal lattice, holding the molecules further apart in the solid than in the liquid — so ice is about 9% less dense than water, floats, and insulates the liquid below. One final discipline: when a molecular substance melts or boils, it is these intermolecular forces that break. The covalent bonds inside the molecules ride through unharmed.
VocabularyKey terms the mark scheme pays for
TrapsMisconceptions that cost marks
ExamWhat examiners want
AQA mark schemes in this section pay for named particles and named forces, in full: 'electrostatic attraction between oppositely charged ions' for ionic, 'shared pair of electrons' for covalent, 'attraction between positive ions and delocalised electrons' for metallic. The unnamed word 'bonds' earns nothing. In any property comparison, identify each substance's structure type first — ionic, metallic, macromolecular or molecular — because the rest of your answer inherits its logic from that classification, and so does the mark scheme.
Shape questions are a fixed routine worth practising to reflex: count the electron pairs (show the arithmetic, adjusting for any charge on an ion), split them into bonding and lone, then state both the shape name and the bond angle — knocking 2.5° off per lone pair. Name shapes from atom positions, never electron positions: water is bent, not tetrahedral. For 6-mark comparisons, structure the answer symmetrically — structure, forces, energy for substance one, then the same triad for substance two — because level-of-response marking rewards completeness on both sides, not brilliance on one.
Two phrasing traps recur. Intermolecular forces act between molecules — 'van der Waals forces between the atoms in a molecule' is a contradiction that costs the mark. And a hydrogen-bonding answer must show both halves of the arrangement: the hydrogen attached to N, O or F, and the lone pair it is attracted to on the neighbouring molecule — a labelled diagram with the hydrogen bond drawn from H to lone pair is often the fastest route to full marks. Section 3.1.3 appears on Papers 1, 2 and 3, and its extended-response questions are where the application weighting (40–45% of the qualification is AO2) shows its teeth: expect unfamiliar molecules, and trust the toolkit.