Chemistry › Bonding and structure
Bonding and structure
Unit C03 · 5 lessons.
Sodium chloride melts at 801 °C and chlorine boils at −34 °C. Graphite conducts and diamond does not, and water boils far above hydrogen sulfide. Each difference comes from bonding and structure: ionic, covalent, dative and metallic bonding, the four crystal structures, the shapes VSEPR predicts, electronegativity and the three kinds of intermolecular force.
The commonest wrong answer is that boiling a simple molecular substance breaks its covalent bonds. It does not. Boiling separates the molecules, so the explanation is about the forces between them, named as hydrogen bonding, permanent dipole–dipole forces or induced dipole–dipole forces (van der Waals forces), and compared in strength.
Questions on it can come up on AQA Papers 1, 2 and 3, OCR A Papers 1, 2 and 3, Edexcel Papers 1, 2 and 3 and Cambridge Papers 1 and 2; the past-paper pages say what else each of those papers covers.
Specification sections: AQA 3.1.3.1, 3.1.3.4, 3.1.3.2, 3.1.3.3, 3.1.3.5, 3.1.3.6, 3.1.3.7 · OCR A 2.2.2, 3.1.1 · Edexcel 2.1 to 2.27 · CAIE 3.2, 4.2, 3.3, 3.4, 3.7, 12.1, 3.5, 13.3, 3.1, 3.6
Notes
Illustrated notes for each topic in this unit, with worked examples and a checklist of objectives.
- Ionic bonding and ionic lattices
Electron transfer and the charges on simple and compound ions, dot-and-cross diagrams for NaCl, MgO and CaCl₂, formula units and counting ions in a mass, the sodium chloride lattice and 6:6 coordination, ion migration and electron density maps as evidence for ions, the strength of ionic bonding and how it depends on ionic charge and radius, ion sizes down a group and across an isoelectronic series, and the melting points, conductivity, brittleness and solubility of ionic compounds.
- Covalent, dative covalent and metallic bonding
Covalent bonds as shared pairs of electrons, dot-and-cross diagrams for H₂O, NH₃, CH₄, HCl, CO₂, O₂ and N₂, lone pairs, double and triple bonds with their lengths and strengths, the unreactivity of nitrogen, H–X bond strength down Group 7, σ and π bonds with an estimate of π bond strength, sp³, sp² and sp hybridisation, molecules outside the octet rule, dative covalent bonds in NH₄⁺, H₃O⁺ and Al₂Cl₆ with the arrow notation, and metallic bonding with its strength, conductivity, malleability and melting points.
- The four crystal structures and their properties
The four crystal structures, ionic, metallic, macromolecular and molecular, with sodium chloride, magnesium, diamond, graphite, silicon dioxide, iodine and ice as examples; the particles and forces in each; the melting point, electrical conductivity and solubility that follow; graphite's layers and conduction; graphene; buckminsterfullerene, C₆₀, as a molecular solid; why ice is less dense than water; enthalpies of fusion and vaporisation and the energy changes at changes of state; drawing the structures; and a full comparison table.
- Shapes of molecules and ions
Electron-pair repulsion theory (VSEPR), the shapes and bond angles for two to six electron pairs (linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral), lone pairs in pyramidal NH₃ and bent H₂O, the ions NH₄⁺, H₃O⁺ and ICl₂⁻, SF₄, T-shaped ClF₃ and square planar XeF₄, why lone pairs go equatorial, the change of shape on forming H₃N→BF₃, double bonds in CO₂, SO₂ and ethene, isoelectronic species such as BH₄⁻, NH₂⁻, CO₃²⁻ and NO₃⁻, ethanol and benzene, sp³, sp² and sp hybridisation, the method for counting electron pairs including ions such as ICl₄⁻, and drawing molecules with wedges and dashes.
- Electronegativity, polarity and intermolecular forces
Electronegativity on the Pauling scale and its trends, polar bonds and partial charges, polar molecules against symmetric molecules such as CO₂ and CCl₄, induced dipole–dipole (London) forces, permanent dipole–dipole forces and hydrogen bonding, the boiling points of the halogens, the alkanes and the hydrides of Groups 4 to 7 with the anomalies of NH₃, H₂O and HF, ice, the surface tension of water, dipole moments, the Cambridge use of van der Waals' forces, and solubility in water, with butan-1-ol against butane, the insolubility of halogenoalkanes and the choice of solvent.
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