Edexcel A-Level Chemistry: Bonding and Structure — Complete Revision Guide (9CH0)
Edexcel A-Level Chemistry: Bonding and Structure — Complete Revision Guide (9CH0)
Bonding is the second major content area of Edexcel A-Level Chemistry (9CH0), and it is the topic that connects atomic structure to everything else. Once you know how electrons are arranged in atoms, the next question is how atoms hold themselves together — and the answer governs almost every property a substance has, from melting point and conductivity to solubility and reactivity. Energetics depends on bond enthalpies. Organic chemistry depends on polarity and bond polarity to predict reaction sites. Inorganic chemistry depends on the bonding type of oxides and chlorides. If you can read a structure off the page and predict how it will behave, the rest of the course becomes far less daunting.
This guide is a topic-by-topic walkthrough of the bonding and structure content in the 9CH0 specification. It covers ionic, covalent, dative and metallic bonding; VSEPR theory and molecular shapes; electronegativity and bond polarity; the three classes of intermolecular forces (London dispersion, permanent dipole-dipole and hydrogen bonding); and the four major structure types (giant ionic, giant covalent, simple molecular, giant metallic) and how each one's properties follow from its structure. For each topic you will find the core ideas, common pitfalls, a worked example or short numerical illustration, and a link into the LearningBro Bonding and Structure course.
What the Edexcel 9CH0 Specification Covers
Edexcel A-Level Chemistry (9CH0) is examined in three written papers. Paper 1 (Inorganic and Physical) is one hour 45 minutes and 90 marks. Paper 2 (Organic and Physical) is the same length and mark allocation. Paper 3 (General and Practical Principles) runs two hours 30 minutes for 120 marks. Bonding sits alongside atomic structure at the start of the specification and is examined directly on Papers 1 and 2, with structure-property reasoning showing up in extended-response questions on Paper 3.
Bonding questions tend to fall into three styles: short recall questions on bond types, drawing tasks on shapes, and extended-response questions linking structure to properties. Together these account for a substantial fraction of every paper. The table below maps the main sub-topics to a typical paper weighting.
| Sub-topic | Spec area | Typical paper weight |
|---|---|---|
| Ionic bonding and ionic structures | Topic 2 | 4-6 marks |
| Covalent and dative bonding | Topic 2 | 4-6 marks |
| Metallic bonding and metals | Topic 2 | 2-4 marks |
| VSEPR molecular shapes | Topic 2 | 4-6 marks |
| Electronegativity and bond polarity | Topic 2 | 3-5 marks |
| Intermolecular forces | Topic 2 | 6-8 marks |
| Structure-property reasoning | Topic 2 | 4-8 marks |
These weights are estimates, modelled on the Edexcel 9CH0 paper format rather than guarantees for any individual paper. What is reliable is that intermolecular forces appear on every paper somewhere, and that a structure-and-properties extended-response question is a near-permanent fixture.
Ionic Bonding and Ionic Structures
Ionic bonding is the electrostatic attraction between oppositely charged ions formed by electron transfer from a metal to a non-metal. The simplest case is sodium chloride: sodium loses its 3s^1 electron to form Na+ with a stable [Ne] configuration; chlorine gains an electron into 3p to form Cl- with a stable [Ar] configuration. The two ions then arrange in a giant lattice in which each cation is surrounded by anions and vice versa.
Ionic bond strength is measured by lattice energy, which depends on two factors: the magnitude of the ionic charges (squared in Coulomb's law) and the inter-ionic distance (the sum of the ionic radii). MgO has a much larger lattice energy than NaCl because Mg^2+ and O^2- carry double charges and are smaller than Na+ and Cl-. The full Born-Haber cycle treatment appears in energetics.
Ionic compounds typically have high melting and boiling points (a great deal of energy is needed to overcome the lattice), are brittle (a layer shift brings like charges together, splitting the crystal), conduct electricity only when molten or dissolved (the ions must be free to move), and are often soluble in polar solvents like water (the ion-dipole interactions with water more than offset the lattice energy).
Worked example. Predict whether MgCl2 or CaCl2 has a higher melting point. Both have +2 cations and Cl- anions, so the difference comes from cation radius. Mg^2+ is smaller than Ca^2+, so the inter-ionic distance is shorter and the lattice energy more negative for MgCl2. MgCl2 has the higher melting point. (In fact, MgCl2 melts at 714 °C and CaCl2 at 772 °C — the actual pattern is influenced by lattice geometry, but the reasoning is the right starting point.)
A common pitfall is to draw the dot-and-cross diagram of an ionic compound with overlapping electron clouds; in ionic bonding the electrons are fully transferred and the diagram should show separate ions with their charges. See the ionic bonding lesson for diagrams of the standard cases.
Covalent Bonding and Dative Bonds
A covalent bond is a shared pair of electrons between two atoms, in which both atoms contribute one electron each. The shared pair sits in a region of overlapping atomic orbitals and is attracted to both nuclei simultaneously. Single, double and triple bonds correspond to one, two and three shared pairs respectively. Double and triple bonds are stronger and shorter than single bonds between the same atoms.
A dative covalent (coordinate) bond is a covalent bond in which both shared electrons come from the same atom. Once formed, a dative bond is identical in every property to an ordinary covalent bond — it is only the origin of the electrons that differs. Standard examples include the bond from N to H+ in NH4+, from N to BF3 in the adduct H3N→BF3, and from H2O to H+ in H3O+. By convention, dative bonds are drawn with an arrow from the lone-pair donor to the acceptor.
Worked example. In the ammonium ion NH4+, three of the N-H bonds form by ordinary covalent bonding (each H contributes one electron, N contributes one). The fourth bond forms by donation of N's lone pair to a proton: N: + H+ → N→H. All four bonds become identical and the ion is tetrahedral.
A common pitfall is trying to draw NH4+ with N-H bond lengths that differ to mark out the dative bond — once formed, all four bonds are equivalent. Another is forgetting that BF3 has only six valence electrons on B, leaving an empty orbital that can accept a lone pair (so BF3 readily forms adducts with NH3, ethers and other donors).
See the covalent bonding lesson for dot-and-cross diagrams and dative bond examples.
Metallic Bonding
Metallic bonding is the electrostatic attraction between a regular lattice of metal cations and a "sea" of delocalised valence electrons. The cations sit at fixed lattice positions; the electrons move freely throughout the structure and are not localised between specific atoms. The strength of metallic bonding depends on the charge of the cation (more charge means stronger attraction to the electron sea) and the size of the cation (smaller means closer approach).
The properties of metals follow directly from this picture. High melting points because a strong electrostatic lattice must be broken (Group 1 metals are exceptions because their +1 charge gives weaker bonding). High electrical conductivity in solid and liquid because the delocalised electrons are free to move. High thermal conductivity by the same mechanism. Malleability and ductility because cations can slide past one another within the electron sea without splitting the structure. Lustre because the delocalised electrons reflect light at metal surfaces.
Worked example. Why is sodium (m.p. 98 °C) much softer and lower-melting than magnesium (m.p. 650 °C)? Sodium is +1 and larger; magnesium is +2 and smaller. Both factors make magnesium's metallic bonding stronger, raising the melting point and the hardness.
A common pitfall is to write "metallic bonds are between cations and anions" — there are no anions in metallic bonding; the second species is the delocalised electron sea. See the metallic bonding lesson for the full picture.
VSEPR and Molecular Shapes
Valence Shell Electron Pair Repulsion (VSEPR) theory says that electron pairs around a central atom repel each other and arrange themselves to be as far apart as possible. Bonding pairs and lone pairs both count, but lone pairs repel slightly more strongly than bonding pairs. The order of repulsion is lone-lone > lone-bonding > bonding-bonding.
| Pairs (total) | Bond pairs / lone pairs | Shape | Bond angle |
|---|---|---|---|
| 2 | 2 / 0 | Linear | 180° |
| 3 | 3 / 0 | Trigonal planar | 120° |
| 4 | 4 / 0 | Tetrahedral | 109.5° |
| 4 | 3 / 1 | Trigonal pyramidal | 107° |
| 4 | 2 / 2 | Bent | 104.5° |
| 5 | 5 / 0 | Trigonal bipyramidal | 90° / 120° |
| 6 | 6 / 0 | Octahedral | 90° |
Worked examples. Methane (CH4) has four bonding pairs around C — tetrahedral, 109.5°. Ammonia (NH3) has three bonding pairs and one lone pair — trigonal pyramidal, 107°. Water (H2O) has two bonding pairs and two lone pairs — bent, 104.5°. The bond angle decreases from CH4 to NH3 to H2O because each lone pair compresses the bonding pairs slightly more.
A common pitfall is to count double bonds as two electron-pair regions when applying VSEPR — they count as one. Another is to forget about lone pairs entirely when predicting shapes. See the VSEPR lesson for shape diagrams and a flowchart for any small molecule.
Electronegativity and Bond Polarity
Electronegativity is the ability of an atom in a covalent bond to attract the bonding electrons. The Pauling scale runs from 0.7 (Cs) to 4.0 (F). When two atoms of equal electronegativity share a bond, the electrons are equally distributed and the bond is non-polar. When the electronegativities differ, the more electronegative atom pulls the bonding pair closer and develops a partial negative charge (δ-), while the other develops a partial positive charge (δ+). The bond is polar.
A polar bond does not always make the molecule polar. Molecular polarity depends on whether the individual bond dipoles cancel by symmetry. CO2 has two polar C=O bonds but is linear, so the dipoles cancel and the molecule is non-polar. H2O has two polar O-H bonds and is bent, so the dipoles add to a net molecular dipole pointing from the O end to the midpoint of the H-H line. CHCl3 is polar; CCl4 is not.
Worked example. Predict whether NH3 and BF3 are polar or non-polar. NH3 has three polar N-H bonds and a lone pair on N; the molecule is trigonal pyramidal, so the bond dipoles do not cancel and there is a net dipole pointing from the H side to the lone pair side. NH3 is polar. BF3 has three polar B-F bonds but is trigonal planar with no lone pair; the dipoles cancel and BF3 is non-polar.
A common pitfall is to label any molecule with polar bonds as polar — molecular geometry must be checked. See the polarity lesson for systematic dipole analysis.
Intermolecular Forces
Intermolecular forces (IMF) are attractions between separate molecules. They are weaker than covalent bonds within a molecule but determine almost every physical property of molecular substances — melting point, boiling point, viscosity, surface tension, solubility. The 9CH0 specification distinguishes three types.
London dispersion forces (also called instantaneous induced dipole-induced dipole, or van der Waals forces) arise because the electrons in any molecule are constantly moving, producing a fluctuating instantaneous dipole. This dipole induces a dipole in a neighbouring molecule, and the two attract. London forces exist between every pair of molecules. They scale with the number of electrons in the molecule (larger molecules have more electrons and stronger dispersion forces) and with the surface area available for contact (linear chains touch over a longer area than spherical isomers).
Permanent dipole-dipole forces act between polar molecules. The δ+ end of one molecule attracts the δ- end of another. They are stronger than dispersion forces of comparable size but weaker than hydrogen bonding.
Hydrogen bonding is a special, much stronger dipole-dipole interaction that occurs when hydrogen is covalently bonded to N, O or F (the "FON" rule). The H-N, H-O or H-F bond is highly polar, and the small bare proton can approach a lone pair on N, O or F in a neighbouring molecule very closely. Hydrogen bonding is responsible for the anomalously high boiling point of water, the lower density of ice than liquid water, the structure of proteins, and the base pairing in DNA.
Worked example. Rank methane (CH4), ethane (C2H6), ethanol (C2H5OH) and ethanoic acid (CH3COOH) by boiling point, lowest first, and explain. CH4: small, non-polar, only weak London forces — boils at -161 °C. C2H6: larger, more London forces — boils at -89 °C. C2H5OH: same size as ethane but with an -OH group, so hydrogen bonding dominates — boils at +78 °C. CH3COOH: hydrogen bonds doubly (each molecule has both an O-H donor and a C=O acceptor, forming dimers) — boils at +118 °C. The trend is dominated by hydrogen bonding strength.
A common pitfall is to call any dipole-dipole attraction "hydrogen bonding" — only N-H, O-H and F-H qualify. Another is to forget London forces when comparing molecules of different sizes. See the intermolecular forces lesson for boiling-point graph practice.
Structure-Property Reasoning
Edexcel's favourite extended-response question on bonding asks you to identify a structure type from properties, or predict properties from structure. There are four canonical structure types you must be able to recognise immediately.
| Structure | Examples | Melting point | Conductivity | Solubility |
|---|---|---|---|---|
| Giant ionic | NaCl, MgO, CaF2 | High | Only molten or aqueous | Often in water |
| Giant covalent | Diamond, graphite, SiO2 | Very high | Only graphite | Insoluble |
| Simple molecular | Iodine, ice, CO2(s) | Low | None | Sometimes (polar in polar) |
| Giant metallic | Na, Mg, Cu, Fe | High (Group 1 lower) | High (solid and liquid) | Insoluble |
Worked example. A solid melts at 2530 °C, does not conduct electricity in any state, and is insoluble in water. Identify the structure type and a likely candidate. Very high melting point and complete non-conductivity rule out ionic (which conducts when molten) and metallic. Insolubility rules out simple molecular for most candidates. The structure is giant covalent. SiO2 (1713 °C) and diamond (3550 °C) are candidates; the value of 2530 °C is consistent with silicon carbide or a similar giant covalent.
A common pitfall is to forget that graphite conducts because of its delocalised electrons between layers. Another is to assume ionic compounds always dissolve in water — many do not (AgCl, CaCO3) when the lattice is too strong relative to the hydration enthalpies. See the structure and properties lesson for a decision tree.
Common Mark-Loss Patterns
- Drawing dot-and-cross diagrams of ionic compounds with overlapping clouds instead of separated ions with charges.
- Calling any δ+/δ- attraction a hydrogen bond — only N-H, O-H and F-H count.
- Stating the shape of a molecule without identifying the lone pairs first.
- Counting a double bond as two electron pair regions in VSEPR (it counts as one).
- Labelling a molecule with polar bonds as polar without checking molecular geometry.
- Forgetting that London forces exist between every pair of molecules and increase with number of electrons.
- Writing "ionic bonds break on melting" — they do not break, the lattice loosens but the bonds remain (the same goes for covalent bonds in simple molecular substances; only the IMF break).
- Confusing the words "bond" and "intermolecular force" — they are different scales and different strengths.
- Predicting properties of metals without using the term "delocalised electrons".
How to Revise This Topic
- Practise dot-and-cross diagrams every day for one week. NaCl, MgCl2, CO2, N2, H2O, NH3, NH4+, H3O+, NO3-, SO4^2-. By the end you should not need to count electrons.
- Learn the VSEPR table by heart, including the angles. Practise switching from molecular formula to shape in under ten seconds.
- Use the boiling-point graph for the hydrides of Groups 4-7 as a single visual story for IMF. The anomalies for H2O, NH3 and HF are the cleanest evidence for hydrogen bonding on the specification.
- Drill the four structure types until you can identify each from a one-line property description. Make a flashcard set.
- Always answer "explain the property" questions in two stages: first describe the structure, then explain how that structure produces the property. Examiners reward this template.
- Use the LearningBro practice quizzes to test under timed conditions.
Linking to Other Topics
Bonding underpins almost every later topic. Energetics uses bond enthalpies and lattice energies — both of which require an understanding of bond strength and structure. Inorganic chemistry tests trends in ionic bonding (Group 2 thermal stability, halide solubility), and the polarity reasoning here applies directly to the metal-aqua complexes in transition-metal chemistry. Organic foundations uses bond polarity to predict reaction sites — the C-Br bond in halogenoalkanes is polar because Br is more electronegative, and that polarity drives nucleophilic substitution. The intermolecular forces section also explains the boiling points of organic homologous series and the solubility of alcohols and amines.
Final Word
Bonding is one of the most generous topics on 9CH0 — the content is finite, the questions are predictable, and a clean understanding pays off across the rest of the course. Drill structures and shapes, learn the trends in intermolecular forces, and practise structure-property extended-response questions until the language flows automatically. The full LearningBro Bonding and Structure course walks through every sub-topic with diagrams, worked examples and AI tutor feedback. Get this section right and the bonding language you build here will support almost every other topic on the specification.