AQA A-Level Chemistry: Bonding and Structure — Complete Revision Guide (7405)
AQA A-Level Chemistry: Bonding and Structure — Complete Revision Guide (7405)
Bonding and structure is the silent backbone of A-Level chemistry. Almost every other topic on AQA 7405 — from energetics and equilibrium through to organic mechanisms and the spectroscopic puzzles on Paper 3 — leans on a small set of ideas first introduced here. Why does sodium chloride melt at 801 °C while chlorine boils at -34 °C? Why is water a liquid at room temperature while methane is a gas? Why does graphite conduct but diamond insulate? Why is the C-Cl bond in chloromethane more reactive than the C-C bond in ethane? Every answer routes back through bonding type, geometry, electronegativity and intermolecular forces.
Get this topic fluent and the rest of the course slots into place. Skim it and you will fight every section. The good news is that bonding is largely conceptual, not calculation-heavy — once the pictures are clear, the marks are reliably easy to win.
Bonding is the most synoptic foundation in 7405. Lattice enthalpy in energetics is a direct measure of ionic-bond strength. The Period 3 chlorides in inorganic chemistry range from ionic NaCl through giant covalent-like behaviour to molecular SiCl4 and PCl5, with melting points that read straight off bonding theory. Polar bonds and dipoles drive every nucleophilic attack and electrophilic addition mechanism in organic foundations. Infrared spectroscopy detects bond polarity directly: only bonds with a changing dipole absorb. Even atomic structure — see atomic structure — sets up the orbital language that bonding then uses.
This guide walks through the bonding content in AQA 7405 topic by topic. It covers ionic bonding and ionic structures; covalent and dative covalent bonding; the shapes of molecules and ions using the VSEPR model; electronegativity and bond polarity; intermolecular forces (London, permanent dipole-dipole, hydrogen bonding); metallic bonding and giant covalent structures; how physical properties reveal bonding type; and the evidence chemists use to determine bonding and structure. For each topic you will find the core concept, the key idea examiners reward, a common pitfall and a link into the LearningBro Bonding and Structure course.
AQA 7405 Specification Coverage
AQA A-Level Chemistry (7405) is examined through Paper 1 (Inorganic and Physical, 2h, 105 marks), Paper 2 (Organic and Physical, 2h, 105 marks) and Paper 3 (any content plus practical skills, 2h, 90 marks). Bonding sits in section 3.1.3 of the specification and is examined across all three papers, most heavily on Paper 1 in short structured items and on Paper 3 as part of synoptic and practical questions.
| Sub-topic | Spec area | Typical paper weight |
|---|---|---|
| Ionic bonding and structures | 3.1.3.1 | 2-4 marks |
| Covalent and dative covalent bonding | 3.1.3.2 | 2-4 marks |
| Shapes of molecules and ions (VSEPR) | 3.1.3.5 | 3-5 marks |
| Electronegativity and bond polarity | 3.1.3.6 | 2-4 marks |
| Intermolecular forces | 3.1.3.7 | 4-6 marks |
| Metallic bonding | 3.1.3.3 | 2-3 marks |
| Giant covalent (macromolecular) structures | 3.1.3.4 | 2-3 marks |
| Properties as evidence of bonding | 3.1.3.4 / 3.1.3.7 | 3-6 marks |
These weights are estimates, modelled on recent 7405 papers. What is reliable is that a VSEPR shape question, an intermolecular forces explanation and a property-from-structure question appear on essentially every series.
Ionic Bonding and Ionic Structures
Ionic bonding is the electrostatic attraction between oppositely charged ions formed by the transfer of electrons from a metal to a non-metal. Sodium loses one electron to form Na+; chlorine gains one to form Cl-. The resulting ions are held in a giant three-dimensional ionic lattice in which each cation is surrounded by anions and vice versa.
The key idea is that ionic bonds are non-directional. The electrostatic force radiates equally in all directions, so ions pack into the most efficient lattice geometry available. In NaCl this is a face-centred cubic arrangement with each ion surrounded by six oppositely charged neighbours. The lattice is the structure — there are no discrete NaCl "molecules".
Ionic-bond strength scales with the product of the ion charges divided by the inter-ionic distance (Coulomb's law). MgO (2+ / 2- ions, small radii) has a much higher melting point than NaCl (1+ / 1- ions, larger radii) for exactly this reason — a pattern that returns when you meet lattice enthalpies in energetics.
A common pitfall is to draw NaCl as a discrete molecule with a single Na-Cl bond. Another is to claim ionic compounds always dissolve in water — many high-charge oxides and carbonates do not, because lattice enthalpy exceeds the hydration enthalpy.
See the ionic bonding lesson.
Covalent Bonding and Dative Covalent Bonding
A covalent bond is a shared pair of electrons between two atoms, where each atom usually contributes one electron. The shared pair occupies a region of high electron density between the nuclei, and the bond is held together by the attraction of both nuclei for that shared density. Covalent bonds are directional — they point along specific internuclear axes, which is why covalent molecules have definite shapes.
A dative (coordinate) covalent bond is a covalent bond in which both electrons come from the same atom — typically a lone pair donated into an empty orbital on the acceptor. Once formed, a dative bond is indistinguishable from any other covalent bond. The classic AQA examples are NH4+ (ammonia donating its lone pair to H+), H3O+ (water doing the same), and AlCl3·NH3 / Al2Cl6 (chlorine lone pairs bridging two aluminium centres).
In displayed formulae a dative bond is drawn as an arrow from donor to acceptor, e.g. H3N→BF3. Examiners accept either an arrow or an ordinary line provided the donation is made clear in the description.
A common pitfall is to assume the four N-H bonds in NH4+ are different lengths or strengths because one is "dative". In fact all four are identical — the dative origin matters for electron counting, not for the final bond. Another is to forget that nitrogen in NH4+ has no lone pair left, which constrains its shape.
See the covalent and dative bonding lesson.
Shapes of Molecules and Ions (VSEPR)
The Valence Shell Electron Pair Repulsion model — usually credited to Gillespie and Nyholm — predicts molecular shape from a simple premise: electron pairs around a central atom repel each other and arrange themselves as far apart as possible. Lone pairs repel more strongly than bonding pairs, compressing bond angles slightly.
To apply VSEPR: count electron pairs (bonding plus lone) around the central atom, assign the parent geometry, then describe the shape using only the atoms.
| Total pairs | Bonding / lone | Parent geometry | Shape | Angle |
|---|---|---|---|---|
| 2 | 2 / 0 | Linear | Linear | 180° |
| 3 | 3 / 0 | Trigonal planar | Trigonal planar | 120° |
| 3 | 2 / 1 | Trigonal planar | Bent | ~117° |
| 4 | 4 / 0 | Tetrahedral | Tetrahedral | 109.5° |
| 4 | 3 / 1 | Tetrahedral | Trigonal pyramidal | ~107° |
| 4 | 2 / 2 | Tetrahedral | Bent | ~104.5° |
| 5 | 5 / 0 | Trigonal bipyramidal | Trigonal bipyramidal | 90° / 120° |
| 6 | 6 / 0 | Octahedral | Octahedral | 90° |
Worked example. SF6 has six bonding pairs and no lone pairs around sulfur, giving an octahedral shape with 90° bond angles. NH3 has three bonding pairs and one lone pair, giving a pyramidal shape with 107° angles — the lone pair pushes the N-H bonds slightly closer together than the 109.5° tetrahedral ideal.
A common pitfall is to forget to count lone pairs when describing shape. Another is to call NH3 "tetrahedral" because the electron geometry is tetrahedral — shape names refer to atoms only.
See the shapes of molecules lesson.
Electronegativity and Bond Polarity
Electronegativity — Pauling's concept — is the ability of an atom in a covalent bond to attract the bonding pair of electrons. It increases across a period (nuclear charge rises, atomic radius falls) and decreases down a group (atomic radius rises, shielding rises). Fluorine is the most electronegative element at 4.0 on the Pauling scale; caesium and francium are the least at around 0.7.
A polar covalent bond forms when the two atoms have different electronegativities. The more electronegative atom carries a partial negative charge (δ-) and the less electronegative atom a partial positive charge (δ+). H-Cl has δ+ on H and δ- on Cl. The larger the electronegativity difference, the more polar the bond. With a difference of about 1.8 or more on Pauling's scale, the bond is best treated as ionic; below about 0.4 it is essentially non-polar.
A molecule is polar overall when its bond dipoles do not cancel by symmetry. CO2 has two polar C=O bonds, but the molecule is linear, so the dipoles cancel — CO2 is non-polar overall. H2O also has two polar bonds, but the bent shape means the dipoles add to give a net dipole — H2O is polar. CCl4 has four polar C-Cl bonds, but the tetrahedral shape gives perfect cancellation. CHCl3 has the same parent shape but the H replaces one Cl, breaking the symmetry — CHCl3 is polar.
A common pitfall is to confuse bond polarity with molecular polarity. Always check the shape before declaring a molecule polar.
See the electronegativity and polarity lesson.
Intermolecular Forces
Intermolecular forces are the attractions between molecules, distinct from the covalent bonds within molecules. They are far weaker than covalent bonds — typically 1-40 kJ mol-1 versus 200-800 kJ mol-1 — but they set the melting point, boiling point, solubility and viscosity of every molecular substance. AQA recognises three types, in order of typical strength.
London (dispersion) forces arise from instantaneous fluctuations in electron density that induce dipoles in neighbouring molecules. They exist between all molecules, polar or not, and grow stronger with molecular size and surface area. The boiling points of the noble gases (He -269 °C, Xe -108 °C) and the alkanes (CH4 -161 °C, C8H18 +126 °C) are determined almost entirely by London forces.
Permanent dipole-dipole forces act between polar molecules. The δ+ end of one molecule attracts the δ- end of another. They typically add a few kJ mol-1 to the London contribution. Butane (non-polar, b.p. -1 °C) and propanone (polar, b.p. 56 °C) have similar molecular masses, but propanone boils 57 °C higher because of dipole-dipole attractions.
Hydrogen bonding is the strongest intermolecular force AQA examines. It occurs when H is bonded directly to N, O or F, and a lone pair on a nearby N, O or F atom can interact with it. It is around ten times stronger than typical dipole-dipole forces. Hydrogen bonding explains the anomalously high boiling point of H2O (vs H2S), the lower density of ice than liquid water, and the structures of DNA and proteins.
A common pitfall is to call London forces "weak van der Waals" and miss that they dominate boiling points across the alkane series. Another is to claim hydrogen bonding exists in HCl (no — Cl is not N, O or F).
See the intermolecular forces lesson.
Metallic Bonding and Giant Covalent Structures
Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a "sea" of delocalised valence electrons. The Drude model of the late nineteenth century captured the essentials and still drives the AQA mark scheme today: the delocalised electrons can move freely, which explains electrical and thermal conductivity; the non-directional electrostatic force lets layers slide past each other without breaking the bonding, explaining malleability and ductility.
Metallic-bond strength rises with the charge on the cation (more electrons donated to the sea) and falls with cation size. Sodium (1+) melts at 98 °C; magnesium (2+) at 650 °C; aluminium (3+) at 660 °C — even though aluminium ions are larger than magnesium ions, the extra charge wins.
Giant covalent structures — or macromolecular structures — have every atom bonded covalently to its neighbours in an extended three-dimensional or two-dimensional network. The whole crystal is essentially one giant molecule. AQA examples include diamond (each C tetrahedrally bonded to four others, extremely hard, electrical insulator), graphite (each C bonded to three others in hexagonal layers with delocalised electrons between layers, soft and conducting), graphene (a single graphite layer, isolated by Geim and Novoselov), silicon (tetrahedral like diamond) and silicon dioxide (SiO4 tetrahedra sharing oxygens).
Giant covalent solids have very high melting points because melting requires breaking strong covalent bonds throughout the lattice, not weak intermolecular forces.
A common pitfall is to describe graphite as "weak covalent bonds between layers". The bonds within layers are strong covalent; the forces between layers are weak London forces.
See the metallic and giant covalent lesson.
Physical Properties and Bonding Types
A property-to-bonding inference is one of AQA's favourite question types. You are given melting point, electrical conductivity in different states and solubility, and asked to deduce the bonding. The reasoning runs in both directions: from bonding to predicted properties, and from observed properties back to bonding.
| Bonding type | Melting point | Conductivity (solid) | Conductivity (molten / aq) | Solubility |
|---|---|---|---|---|
| Ionic | High | None | Conducts when molten or aqueous | Often soluble in water |
| Simple molecular | Low | None | None (not aqueous) | Often soluble in non-polar solvents |
| Giant covalent | Very high | Insulator (except graphite/graphene) | Insulator | Insoluble |
| Metallic | Variable, often high | Good conductor | Good conductor (molten) | Insoluble (reacts with water for reactive metals) |
The reasoning behind each row is consistent with the bonding picture. Ionic solids hold ions in fixed lattice positions so the solid cannot conduct, but melting or dissolving frees the ions to move and carry current. Simple molecular substances melt at low temperatures because only weak intermolecular forces need breaking, and they cannot conduct because there are no charged carriers. Giant covalent solids melt at extreme temperatures because covalent bonds must be broken throughout the lattice. Metals conduct in both solid and liquid states because the electron sea is mobile.
A common pitfall is to give a vague "low melting point" answer for a molecular substance without identifying which intermolecular forces are being overcome. Always name the force.
See the physical properties lesson.
Bonding Evidence and Structure Determination
How do chemists know what bonding actually exists? AQA expects you to recognise the experimental evidence behind the models.
X-ray crystallography measures the positions of atoms and ions in a crystal to within fractions of an angstrom. The lattice geometry of NaCl, the bond lengths and angles in diamond, the layer separation in graphite — all are directly measured. The technique was central to Hodgkin's structural work on penicillin and vitamin B12, and to Franklin's crystallographic data that informed the DNA structure.
Melting and boiling points reveal bonding type through the magnitudes discussed above. The pattern within a series — for example, the boiling points of the hydrogen halides HF, HCl, HBr, HI — exposes which intermolecular forces are at work. HF boils anomalously high because of hydrogen bonding; the others follow a smooth London-forces trend with molecular size.
Electrical conductivity in different states is the cleanest distinction between ionic and covalent solids. An ionic compound is an insulator as a solid but a conductor when molten or aqueous. A molecular solid is an insulator in every state. A metal conducts in solid and liquid forms.
Infrared spectroscopy detects covalent bonds with a changing dipole moment. Polar bonds (C=O, O-H, N-H) absorb strongly; non-polar bonds (C-C, H-H) absorb weakly or not at all. IR is used routinely to confirm the presence of functional groups in organic foundations.
Density and hardness of giant covalent solids and metals reveal the close-packing of atoms and the rigidity of the bonding framework.
A common pitfall is to give "high melting point" as the only evidence for ionic bonding when conductivity-when-molten is the more decisive test. Always combine at least two properties for confident structural assignment.
See the bonding evidence lesson.
Common Mark-Loss Patterns
- Drawing ionic compounds as discrete molecules rather than lattices.
- Calling NH4+ N-H bonds different from each other because one is dative.
- Naming a shape after the electron geometry, including lone pairs.
- Confusing bond polarity with molecular polarity (forgetting to check symmetry).
- Describing London forces as "weak" without explaining how they scale with size.
- Claiming hydrogen bonding exists wherever there is a hydrogen atom.
- Saying graphite has weak covalent bonds (it has strong covalent within layers, weak London between).
- Giving "high melting point" as proof of ionic bonding without conductivity evidence.
- Failing to mention delocalised electrons when explaining metallic conductivity.
- Forgetting that diamond and graphite have the same bonding type but very different properties.
How to Revise This Topic
- Build a VSEPR flashcard set with one card per parent geometry. Drill until naming the shape from a Lewis structure takes under five seconds.
- Memorise representative boiling points for each intermolecular force class. Knowing CH4 (-161 °C), HCl (-85 °C) and H2O (100 °C) by heart anchors every explanation.
- Practise property-to-bonding inferences in batches of ten. Give yourself a property table; identify the bonding type and justify with two pieces of evidence.
- Sketch every Period 3 chloride and predict the bonding from the position in the period. NaCl ionic, MgCl2 ionic, AlCl3 covalent with ionic character, SiCl4 molecular, PCl5 molecular, SCl2 molecular.
- Draw dot-and-cross diagrams for ten species per session, including dative-bonded examples (NH4+, H3O+, AlCl3·NH3).
- Use the LearningBro practice quizzes to test under timed conditions.
Linking to Other Topics
Bonding feeds directly into energetics — lattice enthalpy is the energy released when gaseous ions form an ionic solid, and is the experimental measure of ionic-bond strength. Period 3 trends in inorganic chemistry read directly off the bonding-type table above. Every mechanism in organic foundations — nucleophilic substitution, electrophilic addition, free-radical substitution — begins with an analysis of bond polarity and direction of attack. Even the atomic orbital language from atomic structure returns here when you describe sp3, sp2 and sp hybridisation in advanced organic contexts.
Final Word
Bonding is the conceptual backbone of A-Level chemistry. The arithmetic is light, but the reasoning is rich, and examiners reward clean, structure-led answers. Get fluent with the four bonding types, with VSEPR shapes and angles, with the three intermolecular forces and with how properties reveal structure — and the rest of 7405 becomes a sequence of consequences rather than a list of disconnected facts. The full LearningBro Bonding and Structure course walks through every sub-topic with worked examples, dot-and-cross diagrams, VSEPR builder activities and AI tutor feedback. Lock this section down early and the rest of the course pays you back lesson after lesson.