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VCE Chemistry (Year 11) · Unit 1 — Bonding and Materials · 11m read

Covalent Bonding and Molecular Shapes

Ever wonder why water is a liquid but carbon dioxide is a gas, even though both are made from just three atoms? It all comes down to how atoms share electrons and arrange themselves in space — that's covalent bonding and molecular shape. Understanding these patterns helps you predict how substances behave and why certain molecules are essential for life, from the oxygen you breathe to the proteins in your body.

Learning objective: Describe covalent bonding and predict molecular shapes.

By the end of this lesson you can:

  • I draw Lewis structures.
  • I use VSEPR to predict shapes.
  • I distinguish polar and nonpolar molecules.

Covalent bonding is sharing a snack. Non-metals share electron pairs so both feel full. The shape depends on how electron pairs repel.

Sharing electrons

COVALENT BONDING:
• Between NON-METALS.
• Electrons SHARED to achieve noble gas configuration.

TYPES:
• SINGLE (2 shared electrons): H–H.
• DOUBLE (4 shared): O=O.
• TRIPLE (6 shared): N≡N.

LEWIS (DOT) STRUCTURES:
1. Count total valence electrons.
2. Place central atom (usually least electronegative).
3. Connect atoms with single bonds.
4. Distribute remaining electrons as lone pairs to satisfy octet.
5. Use double/triple bonds if needed.

OCTET RULE:
• Most atoms aim for 8 valence electrons.
• Exceptions: H (2), Be (4), B (6), expanded octet for period 3+.

VSEPR THEORY (Valence Shell Electron Pair Repulsion):
• Electron pairs around central atom repel → adopt geometry that minimises repulsion.
• Shape depends on BONDING and LONE pairs.

COMMON SHAPES:
• 2 bonding pairs, 0 lone → LINEAR (e.g., CO₂). 180°.
• 3 bonding, 0 lone → TRIGONAL PLANAR (BF₃). 120°.
• 4 bonding, 0 lone → TETRAHEDRAL (CH₄). 109.5°.
• 3 bonding, 1 lone → TRIGONAL PYRAMIDAL (NH₃). ~107°.
• 2 bonding, 2 lone → BENT (H₂O). ~104.5°.

POLARITY:
• BOND polarity: from electronegativity difference.
• MOLECULE polarity: depends on shape AND whether polar bonds cancel.
• Example: CO₂ has polar bonds but is nonpolar molecule (linear).
• H₂O bent → polar molecule.

INTERMOLECULAR FORCES (next topic — brief intro):
• Between molecules, weaker than bonds.
• Type depends on polarity.

MOLECULAR SUBSTANCES — PROPERTIES:
• LOW melting / boiling.
• POOR conductors.
• Often don't dissolve in water (except polar molecules).
• Held by weak intermolecular forces.

ISOMERS PREVIEW:
• Same formula, different structure.
• Key in organic chemistry.

COMMON MOLECULES:
• H₂, O₂, N₂ — diatomic.
• CO₂, H₂O, NH₃ — triatomic+ with central atom.

DATIVE (coordinate) BONDS:
• Both electrons from same atom.
• e.g., NH₄⁺ (N donates lone pair to H⁺).

RESONANCE:
• Some molecules have multiple valid Lewis structures.
• Actual structure = average (e.g., O₃, CO₃²⁻).

EXAM TIP:
• Always draw Lewis first, then predict shape.
• Include units / angles in answers.

Covalent essentials

  • Electrons shared: Not transferred.
  • Shape = VSEPR: Repulsion logic.
  • Polarity: Bonds + shape.

Worked example (Medium)

Predict the shape and polarity of water (H₂O).

  1. Lewis. O in centre, 2 H bonds, 2 lone pairs on O.
  2. Electron regions. 4 around O (2 bonds + 2 lone pairs).
  3. Shape. BENT (lone pairs push bonds closer).
  4. Bond polarity. O more electronegative → O–H bonds polar.
  5. Molecule polarity. Bent shape → dipoles don't cancel → POLAR molecule.

Answer: Bent shape (~104.5°); polar molecule.

Worked example (Hard)

Draw the Lewis structure of sulfur dioxide (SO₂), predict its shape and polarity, and explain why the bond angle differs from the ideal trigonal planar angle.

  1. Lewis structure. Total valence electrons: S(6) + 2×O(6) = 18. Central S with two double bonds to O and one lone pair on S. Each O also has two lone pairs.
  2. Electron regions. 3 regions around S: 2 bonding (double bonds) + 1 lone pair. Electron geometry = trigonal planar.
  3. Molecular shape. With one lone pair, the molecular shape is BENT (or V-shaped), similar to water but with a different angle.
  4. Bond angle and polarity. The ideal trigonal planar angle is 120°, but the lone pair repels more strongly than bonding pairs, compressing the O-S-O angle to ~119°. The molecule is POLAR because the bent shape means the two S=O bond dipoles don't cancel.

Answer: SO₂ is bent (~119°) with a lone pair on sulfur. It is a polar molecule because the bond dipoles do not cancel in the asymmetric shape.

Common mistakes

  • Forgetting lone pairs when predicting shape VSEPR counts ALL electron regions (bonding AND lone pairs). NH₃ has 4 electron regions but only 3 bonds — the lone pair makes it pyramidal, not tetrahedral. Fix: Always count lone pairs on the central atom. Molecular shape describes ATOM positions; electron geometry includes lone pairs.
  • Saying CO₂ is polar because it has polar bonds CO₂ has two polar C=O bonds, but its LINEAR shape means the dipoles point in exactly opposite directions and CANCEL. The molecule is nonpolar overall. Fix: Check molecular SHAPE as well as bond polarity. Symmetric shapes can cancel polar bonds.
  • Drawing more bonds than an atom can form Carbon forms 4 bonds, nitrogen 3 (+ 1 lone pair), oxygen 2 (+ 2 lone pairs). Students sometimes draw 5 bonds to carbon. Fix: Count valence electrons carefully. Period 2 elements CANNOT exceed the octet rule.

VSEPR theory — predicting 3D shape from 2D drawings

VSEPR (Valence Shell Electron Pair Repulsion) theory is elegantly simple: electron pairs around a central atom repel each other and arrange themselves to be as far apart as possible, minimising repulsion. This determines molecular geometry.

The key insight is that lone pairs repel more strongly than bonding pairs. The repulsion hierarchy is:

Lone pair–lone pair > Lone pair–bonding pair > Bonding pair–bonding pair

This is because lone pairs are held closer to the nucleus (not shared with another atom) and occupy more angular space. This extra repulsion compresses bond angles:

| Electron regions | Lone pairs | Shape | Ideal angle | Actual angle |
|---|---|---|---|---|
| 4 | 0 | Tetrahedral | 109.5° | 109.5° |
| 4 | 1 | Trigonal pyramidal | 109.5° | ~107° |
| 4 | 2 | Bent | 109.5° | ~104.5° |

Notice how each lone pair reduces the bond angle by about 2-2.5°.

For VCE, you need to recognise five common shapes: linear (180°), trigonal planar (120°), tetrahedral (109.5°), trigonal pyramidal (~107°), and bent (~104.5° or ~119° depending on electron geometry). Drawing these in 3D using wedge-dash notation (solid wedge = toward viewer, dashed line = away) demonstrates understanding and earns credit with examiners.

Understanding molecular shape is not just academic — it explains why enzymes only bind specific substrates (shape complementarity), why drug molecules must have precise 3D structures to work, and why water's bent shape gives it the properties essential for life.

Practice questions

  1. CO₂ is linear and:

    1. Polar.
    2. Nonpolar (bonds cancel).
    3. Ionic.
    4. Unstable.

    (1 marks)

  2. Why is NH₃ polar while BF₃ is nonpolar?

    (3 marks)

Summary

  • Draw Lewis structures.
  • Use VSEPR to predict shapes.
  • Distinguish polar and nonpolar molecules.

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