# Bonding and Structure

> Edexcel International A-Level Chemistry · IAL Chemistry Unit 1 (WCH11)
> Source: https://www.owlsprep.com/study/edexcel-ial-chemistry-u1-bonding-and-structure/

This guide covers all core content for Edexcel IAL Chemistry Unit 1 Topic 3: ionic, covalent and metallic bonding, molecular shape prediction, bond and molecular polarity, and structure-property links.

**Prerequisites:** [Edexcel IAL Chemistry Unit 1 Atomic Structure](https://www.owlsprep.com/study/edexcel-ial-chemistry-u1-atomic-structure/); [Periodic Table Group Trends](https://www.owlsprep.com/study/edexcel-ial-chemistry-u1-periodic-trends/)

## Learning objectives

- Explain ionic bond formation, lattice structure and factors affecting ionic bond strength
- Draw dot-and-cross diagrams for ionic, covalent and dative covalent species
- Predict bond polarity and molecular polarity using electronegativity and geometry
- Apply electron-pair repulsion theory to predict shapes and bond angles of simple molecules
- Describe metallic bonding and link structure to physical properties of metals

## Ionic Bonding and Lattices

**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.

> **warning**
>
> Polarisation occurs when a small, highly charged cation (high polarising power) distorts the electron cloud of a large anion (high polarisability). High polarisation leads to partial covalent character in ionic bonds, e.g. in AgCl.

**Worked example:** 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.

## Covalent Bonding and Polarity

**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.

**Worked example:** 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.

## Molecular Shapes (Electron-Pair Repulsion Theory)

**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.

**Worked example:** 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.

## Metallic Bonding

**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.

**Worked example:** 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).

## Common pitfalls

- **Wrong:** Drawing shared electrons in dot-and-cross diagrams for ionic compounds
  - Why it fails: Ionic bonds involve full electron transfer, not sharing, so electrons are completely transferred to the non-metal ion
  - Correct: Show the metal ion with no outer valence electrons, the non-metal ion with full outer shell, and clearly label ion charges
- **Wrong:** Stating that all molecules with polar bonds are polar
  - Why it fails: Polar bonds cancel out if the molecule has symmetric geometry, leading to no net dipole
  - Correct: Always consider both bond polarity and molecular geometry to determine if a molecule has a net dipole
- **Wrong:** Using hybridisation or AXE notation to explain molecular shapes
  - Why it fails: These concepts are out of scope for Edexcel IAL Unit 1, and marks are only awarded for explanations using electron-pair repulsion theory
  - Correct: 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:** Forgetting to account for ion charge when counting valence electrons for charged species
  - Why it fails: Incorrect electron counts lead to wrong dot-and-cross diagrams and lost marks
  - Correct: Add electrons for negative charges, subtract electrons for positive charges when calculating total valence electrons for a species
- **Wrong:** Claiming metallic conductivity comes from moving ions
  - Why it fails: In solid metals, cations are fixed in the lattice, only delocalised electrons are mobile
  - Correct: Explain that electrical conductivity in metals is due to movement of delocalised valence electrons through the lattice

## 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 |

## 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.

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