amazon.co.ukBonding and Structure
A* revision for ionic, covalent and metallic bonding, electronegativity, intermolecular forces and VSEPR shapes, with exam-style questions for AQA, Edexcel, OCR, WJEC and CCEA A-Level Chemistry.
🎯 Learning Objectives
- Draw dot-and-cross diagrams for ionic, covalent and dative bonded species
- Explain properties of ionic, metallic, giant covalent and simple molecular substances from their structure and bonding
- Use electronegativity differences to predict bond polarity and whether a molecule is polar overall
- Describe London forces, dipole-dipole forces and hydrogen bonding, and predict which occur between given molecules
- Predict shapes and bond angles using VSEPR, explaining deviations caused by lone pairs
- Compare melting points, boiling points and conductivity of substances using the forces that must be overcome
📋 Key Definitions and Core Concepts
Ionic bonding: The electrostatic attraction between oppositely charged ions formed by electron transfer from a metal to a non-metal.
Covalent bond: The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
Dative (coordinate) bond: A covalent bond in which both electrons of the shared pair come from the same atom.
Metallic bond: The electrostatic attraction between positive metal ions and delocalised electrons.
Electronegativity: The ability of an atom to attract the pair of electrons in a covalent bond (Pauling scale).
Intermolecular force: An attractive force between molecules: London (induced dipole-dipole), permanent dipole-dipole, or hydrogen bonding.
Hydrogen bond: A strong dipole-dipole attraction between an electron-deficient H bonded to N, O or F and a lone pair on a neighbouring N, O or F.
Delocalised electrons: Electrons not associated with any one atom or bond, free to move through the structure (metals, graphite).
Giant covalent lattice: A structure of atoms bonded covalently throughout, e.g. diamond, graphite, SiO₂.
Ionic Bonding, Giant Lattices and Properties
Ionic bonding occurs between a metal and a non-metal: electrons are transferred from the metal to the non-metal, forming positive and negative ions held together by strong electrostatic forces of attraction in all directions. Every ion attracts oppositely charged neighbours, producing a giant ionic lattice — the diagram convention is a regular 3D array of alternating cations and anions. A dot-and-cross diagram shows outer-shell electron transfer: the metal's electron(s) are drawn as dots moved onto the non-metal's crosses.
Lattice strength depends on ionic charge and ionic radius: MgO (Mg²⁺, O²⁻) melts at 2852 °C versus NaCl (Na⁺, Cl⁻) at 801 °C because ±2 charges and smaller ions give a higher charge density and much stronger electrostatic attraction. This explains trends in melting points and is the same reasoning used for lattice enthalpies (see Thermodynamics and Born-Haber cycles).
Properties of ionic compounds follow directly from the lattice: (1) high melting and boiling points because strong electrostatic forces must be overcome throughout the lattice; (2) conduction only when molten or dissolved — then ions are mobile and carry charge, whereas in the solid the ions are fixed in place; (3) solubility in polar solvents like water, which hydrate and stabilise the ions. Exam answers must state the force broken on melting: it is electrostatic attraction between ions, not covalent bonds, and it is attraction not repulsion that is overcome.
Covalent, Dative and Metallic Bonding
Covalent bonding is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms — found between non-metals. A single bond shares one pair, a double bond two pairs, a triple bond three; bond enthalpy rises and bond length falls in the order single > double > triple. In dot-and-cross diagrams shared pairs overlap the two atoms; outer shells only are drawn, and you should state that both nuclei attract the shared pair.
A dative (coordinate) covalent bond forms when both electrons in the shared pair come from the same atom: the donor must have a lone pair and the acceptor an empty orbital. Examples: NH₄⁺ (N donates its lone pair to H⁺), H₃O⁺, the [Ag(NH₃)₂]⁺ complex, CO (one normal C–O plus one dative). Dative bonds behave exactly like ordinary covalent bonds once formed — same length, same strength.
Metallic bonding is the attraction between positive metal ions and delocalised electrons — a lattice of cations in a sea of mobile electrons. Delocalisation explains properties: high melting points (strong attraction, rising with ion charge — Mg melts higher than Na because Mg²⁺ attracts two electrons per ion), electrical and thermal conductivity (mobile electrons carry charge and heat), and malleability/ductility (layers of ions can slide without breaking the bonding because delocalised electrons hold the structure together in any arrangement).
Simple molecular substances (H₂, H₂O, CO₂, NH₃, I₂) contain strong covalent bonds within molecules but weak intermolecular forces between them: low melting and boiling points (only IMFs overcome), poor conductivity (no mobile ions or electrons), and solubility varies with polarity. This within/between distinction earns marks constantly — melting a molecular solid breaks intermolecular forces, never covalent bonds.
Electronegativity, Polarity and Intermolecular Forces
Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond, measured on the Pauling scale. It increases across a period (nuclear charge rises, radius falls) and decreases down a group (radius and shielding rise). Fluorine is the most electronegative element (4.0), then O (3.5), N and Cl (3.0).
If two bonded atoms have similar electronegativity the bond is non-polar covalent; a large difference gives ionic character (rough guide: difference > 1.7 ionic, < 0.4 effectively non-polar, between → polar covalent). In a polar bond the more electronegative atom carries a partial charge δ− and the other δ+; the bond has a permanent dipole. Whether the whole molecule is polar depends on shape: symmetrical molecules (CO₂, CCl₄, BF₃, CH₄) have dipoles that cancel, giving non-polar molecules despite polar bonds; asymmetrical ones (H₂O, NH₃, CHCl₃) are polar overall.
Three intermolecular forces (weakest to strongest): (1) London (induced dipole-dipole) forces — exist between all molecules; caused by instantaneous fluctuations in electron distribution inducing dipoles in neighbours; strength increases with number of electrons (higher Mr) and with surface contact (long straight-chain alkanes boil higher than branched isomers). (2) Permanent dipole-dipole forces — between polar molecules only, in addition to London forces. (3) Hydrogen bonding — the strongest IMF, occurring when H is bonded directly to N, O or F (small, highly electronegative atoms with lone pairs); shown as a dashed line from the H to the lone pair of N/O/F on a neighbouring molecule.
Consequences to quote: H₂O boils at 100 °C far above H₂S (−60 °C) despite H₂S having more electrons, because water hydrogen-bonds; ice is less dense than water because H-bonds hold molecules in an open lattice; ammonia's high solubility in water involves H-bonding. Exam rule: state the exact IMF, name the species it acts between, and link strength to energy needed to separate molecules.
Ordering IMFs in comparison answers: every pair of molecules has London forces; polar molecules add dipole-dipole; H-bonding is a special case of dipole-dipole requiring H on N, O or F. When ranking boiling points, first check molecular size (London forces grow with electrons), then polarity, then H-bonding — e.g. propane (Mr 44) vs ethanal (Mr 44): ethanal boils higher because permanent dipole-dipole adds to similar London forces; ethanol (Mr 46) boils far higher still because of H-bonding. Solubility follows the same logic: polar/H-bonding solutes dissolve in polar solvents ('like dissolves like') — ethanol miscible with water (H-bonds between molecules), hexane immiscible (London forces only, cannot disrupt water's H-bond network energetically). Non-polar molecules dissolve in non-polar solvents. These comparisons are standard 3-4 markers: name the IMFs on both sides, then link strength to energy.
Shapes of Molecules: VSEPR and Bond Angles
VSEPR theory: electron pairs around a central atom repel each other and arrange to be as far apart as possible. Repulsion strength: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair — a lone pair is held closer to the central atom, so it repels more strongly and compresses bond angles by about 2.5° per lone pair. A double bond counts as one bonding region.
The shapes to know, with electron-pair counting: 2 regions → linear, 180° (CO₂, BeCl₂); 3 → trigonal planar, 120° (BF₃, SO₃, CO₃²⁻); 4 → tetrahedral, 109.5° (CH₄, NH₄⁺, SO₄²⁻); 5 → trigonal bipyramidal, 90° and 120° (PCl₅); 6 → octahedral, 90° (SF₆). Lone-pair variants: NH₃ (3 bp + 1 lp) trigonal pyramidal 107°; H₂O (2 bp + 2 lp) non-linear/bent 104.5°. Exam answers must name the shape, give the angle, and explain the deviation by repulsion order.
Method for any unfamiliar molecule: (1) count bonding regions around the central atom; (2) count lone pairs on the central atom; (3) pick the base shape from total regions; (4) subtract 2–2.5° per lone pair for the exact angle. For ions, add one electron for each negative charge (NH₄⁺ has no lone pair on N, but OH₃⁺/H₃O⁺ keeps one) and subtract one for each positive charge before counting lone pairs.
Structure and Bonding: Linking to Physical Properties
Four structure types anchor the comparison. Giant ionic (NaCl, MgO): high MP/BP; conducts when molten or aqueous; often soluble in water. Giant covalent (diamond, silicon dioxide, graphite): very high MP/BP — strong covalent bonds must be broken; insoluble; conductivity depends on delocalisation (diamond none, graphite yes in one plane because each C is bonded to only three others leaving one delocalised electron per atom). Giant metallic: high MP/BP, conducts as solid and liquid, malleable. Simple molecular: low MP/BP, non-conducting, H-bonded ones dissolve in water.
Model comparisons: diamond vs graphite — both giant covalent, both high MP because covalent bonds must break to melt, but graphite conducts (delocalised electrons between layers, which can also slide giving lubricating softness) while diamond does not (all four electrons in σ bonds). MgO vs NaCl — higher charges, smaller ions, stronger lattice. HF vs HCl — H-bonding vs dipole-dipole/London only. I₂(s) vs NaCl(s) — weak London forces between I₂ molecules vs strong electrostatic forces in the lattice.
Graphene and fullerenes are sometimes asked about: graphene is a single layer of graphite with delocalised electrons making it conductive and extremely strong; fullerenes are simple molecular (C₆₀), so low melting points despite covalent bonds within each sphere. Structure–property questions are marked by the causal chain: structure type → which forces/bonds must be overcome → energy required → property observed. Always name the force, never just say 'strong bonds'.
Exam technique for structure-property questions: the mark scheme is built from a causal chain — name the structure type, name the exact force or bond broken on melting/boiling, link to the energy required, conclude the property. 'Strong bonds' scores nothing: it must be 'strong electrostatic attraction between oppositely charged ions' or 'strong covalent bonds throughout the lattice' or 'weak London forces between molecules'. Conductivity answers must identify the mobile charge carrier (delocalised electrons in metals/graphite; mobile ions when molten/dissolved) and state why it is immobile in the other state. When comparing two substances, answer both halves: why A is high AND why B is low — half-comparisons lose half the marks. Drawing conventions: dot-and-cross diagrams show outer shells only with bonding pairs overlapping; displayed formulae show every bond; skeletal formulae omit C and H labels on the chain. Fluency converting between the three is assumed at A-level and tested indirectly throughout organic chemistry.
💡 Worked Example
Exam-style: Explain why H₂O has a much higher boiling point than H₂S despite H₂S having a higher relative molecular mass.
Mark-scheme logic: H₂O molecules form hydrogen bonds: O is small and highly electronegative, leaving the H strongly δ+ and exposing lone pairs, so each molecule attracts neighbouring molecules strongly. H₂S has only permanent dipole-dipole and London forces (S is larger and less electronegative). Hydrogen bonds are much stronger than the other IMFs, so more energy is needed to separate the molecules — hence water boils at 100 °C while H₂S boils at −60 °C. Mark points: name hydrogen bonding; conditions (H on O, lone pair); compare with H₂S IMFs; energy-to-separate link. (AQA 3.1.3.10-3.1.3.11)
❓ Practice Questions & Mark Schemes
Predict the shapes and bond angles of NH₃, CO₂ and SF₆.
Answer: NH₃: 3 bp + 1 lp → trigonal pyramidal, 107° (lp repels more than bp, compressing from 109.5°). CO₂: 2 double-bond regions → linear, 180°. SF₆: 6 bp → octahedral, 90°.
Rank CH₄, NH₃, H₂O and HF by boiling point and explain the order.
Answer: CH₄ < NH₃ < HF < H₂O. CH₄ lowest: London forces only. NH₃, HF, H₂O all hydrogen-bond; strength of H-bonding depends on electronegativity of N, F, O — H-F bonds are individually strongest, but water forms more H-bonds per molecule on average (2 donors + 2 acceptors), giving H₂O the highest boiling point.
Why does MgO have a higher melting point than NaCl?
Answer: Mg²⁺ and O²⁻ carry ±2 charges (vs ±1) and are smaller ions, giving higher charge density. Electrostatic attraction between oppositely charged ions in the giant ionic lattice is therefore much stronger in MgO, so more energy is needed to overcome it — 2852 °C vs 801 °C.
Explain why graphite conducts electricity but diamond does not.
Answer: Graphite: each carbon bonds to only three others, so one electron per atom is delocalised between the layers and mobile — carries charge. Diamond: all four outer electrons are localised in σ covalent bonds, no mobile charge carriers, so no conduction.
Explain why the boiling point of butane is higher than that of 2-methylpropane (same M_r).
Answer: Both have London forces only. Butane's straight chain has greater surface contact between molecules than branched 2-methylpropane, so stronger induced dipole-dipole attractions — more energy needed to separate the molecules.
Draw and describe the bonding in NH₄⁺, including the dative bond.
Answer: N forms three normal N–H covalent bonds with H atoms, then donates its lone pair to H⁺ forming a fourth (dative) bond. 4 bonding regions, no lone pairs → tetrahedral, 109.5°. All four bonds are identical in length and strength once formed.