Unit Overview
Structure 3: Chemical Bonding
IB Chemistry SL· 5 min read 📊 15-20% of overall IB Chemistry SL exam
1. Unit at a Glance
This unit builds on the atomic structure you learned in Unit 2, moving from individual atoms to how they interact to form the compounds and materials around us. We progress from the two extreme bonding types (ionic and covalent) to the geometry of covalent molecules, forces between molecules, metallic bonding, and finally orbital hybridization.
A core theme across all sub-topics is linking microscopic bonding structure to macroscopic properties. You will consistently connect how atoms are bonded to observable properties like boiling point, electrical conductivity and solubility, which are heavily assessed in IB exams.
This unit is divided into 6 core sub-topics:
Ionic bonding and structure
Learn how ionic bonds form, describe ionic lattice structure, and relate structure to ionic compound properties.
★★⏱ 15 min
Covalent bonding
Covers covalent bond formation, bond polarity, bond energy, and properties of simple and giant covalent compounds.
★★⏱ 15 min
Shapes of molecules and ions
Apply VSEPR theory to predict the shape, bond angle and overall polarity of molecules and polyatomic ions.
★★★⏱ 20 min
Intermolecular forces
Learn the four main types of intermolecular forces and their impact on melting/boiling points and solubility.
★★★⏱ 20 min
Metallic bonding
Understand metallic bonding structure and how it explains the properties of metals and alloys.
★★⏱ 10 min
Hybridization of atomic orbitals
Identify sp, sp² and sp³ hybridization and relate orbital mixing to observed molecular geometry.
★★★★⏱ 15 min
2. Common Pitfalls
Wrong move:
Confusing intermolecular forces with intramolecular covalent/ionic bonds
Why:
Intermolecular forces are between molecules, much weaker than bonds within molecules, but are the cause of boiling/melting point trends
Correct move:
Always clarify if you are describing forces within a molecule or between molecules when explaining bulk properties
Wrong move:
Assuming all molecules with polar bonds are polar overall
Why:
Symmetrical molecular shape cancels out bond dipoles, resulting in non-polar molecules even with polar bonds
Correct move:
Always check molecular shape and symmetry after identifying polar bonds to determine overall molecular polarity
Wrong move:
Ignoring lone pair repulsion in VSEPR calculations
Why:
Lone pairs repel more strongly than bonding pairs, which reduces bond angles from ideal values
Correct move:
Always count all electron domains (lone pairs + bonding pairs) and adjust for greater lone pair repulsion
3. Quick Reference Cheatsheet
Concept | Key Information |
|---|---|
VSEPR Electron Domains | 2 = linear (180°), 3 = trigonal planar (120°), 4 = tetrahedral (109.5°) |
Intermolecular Force Strength | Ion-dipole > Hydrogen bonding > Dipole-dipole > London dispersion forces |
Bond Polarity Rule | Approx ΔEN > 1.8 = ionic, 0.4 < ΔEN < 1.8 = polar covalent, ΔEN < 0.4 = non-polar covalent |
Hybridization Mapping | 2 electron domains = sp, 3 = sp², 4 = sp³ |
Ionic Properties | High melting point, conducts electricity only when molten/dissolved, brittle |
Metallic Bonding Definition | Electrostatic attraction between positive metal cations and a delocalized sea of electrons |
What's Next
Begin your study of this unit with the first sub-topic on ionic bonding, to build a foundational understanding of how different bonding types form and their properties. Once you complete all 6 sub-topics in this bonding unit, you can move on to the next unit on energetics, where you will connect bonding structure to energy changes in chemical reactions.