Bonding and Structure
Edexcel International A-Level Chemistry· Spec points 3.1–3.22 (Topics 3A–3D)· 45 min read
1. Ionic Bonding and Lattices★★☆☆☆⏱ 10 min
Ionic Bond
The net electrostatic force of attraction between oppositely charged ions in a giant ionic lattice.
Example:
Bonding between Na⁺ and Cl⁻ ions in sodium chloride
Ionic bonds form when metal atoms lose electrons to form positive cations, and non-metal atoms gain electrons to form negative anions. Evidence for ionic structure includes high melting/boiling points, conductivity when molten or dissolved, electron density maps showing zero electron density between ions, and ion migration during electrolysis.
Ionic bond strength depends on ion charge (higher charge = stronger attraction) and ionic radius (smaller radius = stronger attraction). Ionic radii increase down a group as more electron shells are added, and decrease across isoelectronic series (e.g. N³⁻ → Al³⁺) as nuclear charge increases with constant electron count.
Draw the dot-and-cross diagram for magnesium oxide, and explain why it has a higher melting point than sodium chloride.
- 1
Magnesium (group 2) loses 2 valence electrons to form Mg²⁺; oxygen (group 16) gains 2 electrons to form O²⁻.
- 2
Dot-and-cross diagram: Show Mg with no outer electrons and 2+ charge, O with 8 outer valence electrons (6 from O, 2 from Mg) and 2- charge.
- 3
MgO has higher ion charge (+2/-2 vs +1/-1 for NaCl) and smaller ionic radii, so electrostatic attractions are stronger, requiring more energy to overcome, leading to a higher melting point.
Exam tip:
Always label ion charges on ionic dot-and-cross diagrams, and explicitly link ion charge and radius to bond strength when explaining melting points.
2. Covalent Bonding and Polarity★★★☆☆⏱ 12 min
Covalent Bond
Electrostatic attraction between a shared pair of electrons and the positively charged nuclei of two bonded atoms. Dative covalent bonds form when one atom donates both electrons in the shared pair.
Example:
Dative bond in NH₄⁺ forms when the lone pair on N is donated to an H⁺ ion
Dot-and-cross diagrams represent single, double, triple and dative covalent bonds, including for species like Al₂Cl₆ and NH₄⁺. Giant covalent structures (diamond, graphite, graphene) have very high melting points as strong covalent bonds extend across the entire lattice: diamond is used in cutting tools, graphite in electrodes, and graphene in high-strength electronics.
The larger the electronegativity difference between two bonded atoms, the more polar the bond, with a continuous spectrum between fully ionic and fully covalent bonds. Polar bonds create partial charges, but a molecule is only polar if dipoles from polar bonds do not cancel due to symmetric geometry.
Explain why carbon dioxide (CO₂) is non-polar even though it contains polar C=O bonds.
- 1
Oxygen is more electronegative than carbon, so each C=O bond is polar, with a dipole pointing towards the oxygen atom.
- 2
CO₂ has a linear shape, so the two dipoles point in opposite directions and cancel completely.
- 3
There is no net dipole moment, so CO₂ is non-polar.
Exam tip:
When explaining molecular polarity, you must state both the polarity of individual bonds AND the molecule's geometry to show whether dipoles cancel out.
3. Molecular Shapes (Electron-Pair Repulsion Theory)★★★★☆⏱ 15 min
Electron-Pair Repulsion Theory
Electron pairs (bonding and lone pairs) around a central atom repel each other and arrange as far apart as possible to minimise repulsion. Lone pairs repel ~2.5° more strongly than bonding pairs, reducing bond angles.
Example:
H₂O has 2 bonding pairs and 2 lone pairs, so it has a bent shape with bond angle 104.5°
Key shapes and bond angles to recall: BeCl₂ (linear, 180°), BCl₃ (trigonal planar, 120°), CH₄ (tetrahedral, 109.5°), NH₃ (pyramidal, 107°), NH₄⁺ (tetrahedral, 109.5°), H₂O (bent, 104.5°), CO₂ (linear, 180°), gaseous PCl₅ (trigonal bipyramidal, 90°/120°), SF₆ (octahedral, 90°), C₂H₄ (planar, 120° per C). Apply the theory to analogous species by counting central atom electron pairs.
Predict the shape and bond angle of the H₃O⁺ ion, explaining your reasoning.
- 1
Central O has 6 valence electrons, subtract 1 for the positive charge, 3 used in bonding to H atoms, leaving 1 lone pair of electrons.
- 2
There are 3 bonding pairs and 1 lone pair around O. Lone pairs repel more strongly than bonding pairs, so the shape is pyramidal.
- 3
The bond angle is ~107°, slightly less than the tetrahedral 109.5° due to extra repulsion from the lone pair.
Exam tip:
Do not use hybridisation or AXE notation in answers. Explicitly count bonding and lone pairs, and reference the stronger repulsion from lone pairs to justify bond angles.
4. Metallic Bonding★★☆☆☆⏱ 8 min
Metallic Bond
Net electrostatic attraction between a lattice of positive metal cations and a delocalised sea of mobile valence electrons.
Example:
Bonding in magnesium: attraction between Mg²⁺ cations and delocalised outer electrons
Metals have high melting/boiling points due to strong electrostatic attractions between cations and delocalised electrons: higher ion charge and smaller ionic radius increase bond strength and melting point. Metals conduct electricity in solid and molten states as delocalised electrons are free to move and carry charge.
Explain why magnesium has a higher melting point than sodium.
- 1
Magnesium forms Mg²⁺ ions while sodium forms Na⁺ ions, so Mg ions have higher charge and smaller ionic radius.
- 2
Magnesium has more delocalised electrons per ion than sodium.
- 3
Electrostatic attractions in Mg are stronger, so more energy is required to overcome them, leading to a higher melting point.
Exam tip:
When explaining metallic conductivity, reference mobile delocalised electrons, not moving ions (cations are fixed in the lattice in solid state).
5. Common Pitfalls
Wrong move:
Drawing shared electrons in dot-and-cross diagrams for ionic compounds
Why:
Ionic bonds involve full electron transfer, not sharing, so electrons are completely transferred to the non-metal ion
Correct move:
Show the metal ion with no outer valence electrons, the non-metal ion with full outer shell, and clearly label ion charges
Wrong move:
Stating that all molecules with polar bonds are polar
Why:
Polar bonds cancel out if the molecule has symmetric geometry, leading to no net dipole
Correct move:
Always consider both bond polarity and molecular geometry to determine if a molecule has a net dipole
Wrong move:
Using hybridisation or AXE notation to explain molecular shapes
Why:
These concepts are out of scope for Edexcel IAL Unit 1, and marks are only awarded for explanations using electron-pair repulsion theory
Correct move:
Count bonding pairs and lone pairs around the central atom, explain that lone pairs repel more strongly than bonding pairs, and link to shape and bond angle
Wrong move:
Forgetting to account for ion charge when counting valence electrons for charged species
Why:
Incorrect electron counts lead to wrong dot-and-cross diagrams and lost marks
Correct move:
Add electrons for negative charges, subtract electrons for positive charges when calculating total valence electrons for a species
Wrong move:
Claiming metallic conductivity comes from moving ions
Why:
In solid metals, cations are fixed in the lattice, only delocalised electrons are mobile
Correct move:
Explain that electrical conductivity in metals is due to movement of delocalised valence electrons through the lattice
6. Quick Reference Cheatsheet
Bond/Structure Type | Key Features | Typical Properties | Common Examples |
|---|---|---|---|
Ionic | Giant lattice of oppositely charged ions, electrostatic attraction | High melting/boiling point, conducts when molten/dissolved, brittle | NaCl, MgO, CaF₂ |
Simple Covalent | Discrete molecules with shared electron pairs | Low melting/boiling point, non-conductive, often insoluble in water | H₂O, CO₂, CH₄ |
Giant Covalent | Continuous network of strong covalent bonds | Very high melting point, mostly non-conductive | Diamond, graphite, graphene, SiO₂ |
Metallic | Lattice of cations in delocalised electron sea | High melting point, conducts solid/molten, malleable | Fe, Mg, Cu, alloys |
Key Molecular Shapes | Electron pair repulsion, lone pairs reduce angle by ~2.5° | Linear (180°), trigonal planar (120°), tetrahedral (109.5°), pyramidal (107°), bent (104.5°) | BeCl₂, BCl₃, CH₄, NH₃, H₂O |
7. Frequently Asked
Do I need to remember electronegativity values for the exam?
No, Pauling electronegativity values are provided in your official exam data booklet. You only need to use given values to calculate differences and predict bond polarity.
Can I use AXE notation for VSEPR shape questions?
Edexcel does not accept AXE notation for Unit 1 answers. You must explicitly count bonding pairs and lone pairs, and use electron-pair repulsion theory to justify shapes and bond angles.
What's Next
Now that you have mastered bonding and structure, you can build on this knowledge to understand intermolecular forces (Unit 1 Topic 7), which explain the physical properties of simple covalent substances, and later lattice energy and Born-Haber cycles in Unit 2 Topic 12. This content is also foundational for organic reaction mechanisms and group 2/7 periodic trends, which appear frequently in Unit 1 multiple choice and structured questions. Practice past paper questions on dot-and-cross diagrams, molecular shape prediction and polarity to reinforce your understanding and identify knowledge gaps.