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Q: What does H2 Chemistry Notes: Topic 2 - Chemical Bonding cover? A: Consolidate VSEPR shapes, bond polarity, intermolecular forces, and lattice energy trends for Core Idea 2 (Chemical Bonding) in the 2026 H2 Chemistry syllabus.
Chemical bonding connects atomic structure to macroscopic behaviour: lattice energies govern melting points, molecular geometry decides intermolecular forces, and bond polarity predicts reaction pathways. This note organises the skills required for Paper 2 and Paper 3 questions, supported by exam-style worked examples. For the full set of H2 Chemistry revision resources, refer to https://eclatinstitute.sg/blog/h2-chemistry-notes. For the full topic map and paper weightings, see our H2 Chemistry Syllabus 2026-27 overview.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-01-13. Core Idea 2 expectations remain assessed across Papers 1–3.
Quick revision box
What this topic tests: Shapes, polarity, intermolecular forces, and structure-property links.
Top mistakes to avoid: Confusing molecular shape with electron geometry; listing IMF without comparing strength; vague melting-point explanations.
20-minute sprint plan: 5 min VSEPR recap; 10 min polarity/IMF compare questions; 5 min structure-property summary.
1 Types of Chemical Bonds
Bonding model
Key features
Typical exam focus
Ionic
Electrostatic attraction between cations and anions in a lattice.
Explaining trends in melting point, solubility, conductivity.
Covalent
Shared pair(s) of electrons between atoms.
Bond polarity, molecular shape, sigma vs pi bonds.
Lattice of positive ions in a sea of delocalised electrons.
Conductivity, malleability, strength across the periodic table.
1.1 Lattice Energy Trends
Lattice energy (lattice enthalpy of formation) measures enthalpy change when one mole of ionic solid forms from gaseous ions. A simplified Born-Landé equation highlights core dependencies:
The lattice energy U is roughly proportional to (z+×z−) divided by r0, where z₊ and z₋ are the ionic charges and r₀ is the internuclear distance. Larger charge products and shorter separations make lattice formation more exothermic (or lattice dissociation more endothermic), raising melting points and hardness.
1.2 Covalent Bond Strength and Length
Stronger bonds have higher bond dissociation enthalpies and shorter lengths. The syllabus expects students to relate strength to orbital overlap: sigma bonds involve head-on overlap (stronger), whereas pi bonds involve side-on overlap (weaker). Recall that double bonds comprise one sigma and one pi bond.
When numerical comparisons are needed, use the bond enthalpies tabulated in the SEAB Chemistry Data Booklet (exams from 2026) instead of memorised values.
2 Molecular Geometry and Polarity
2.1 VSEPR Workflow
Determine the central atom's steric number: count bonded pairs + lone pairs.
Assign electron-domain geometry (e.g. tetrahedral for steric number 4).
Deduce molecular shape by considering only bonding pairs.
Judge polarity: combine vector sum of bond dipoles; consider symmetry.
Steric number
Electron-domain geometry
Common shapes
Bond angles
3
Trigonal planar
Trigonal planar, bent
120∘, <120∘
4
Tetrahedral
Tetrahedral, trigonal pyramidal, bent
109.5∘,107∘,104.5∘
5
Trigonal bipyramidal
Seesaw, T-shaped, linear
120∘ and 90∘ combinations
6
Octahedral
Square pyramidal, square planar
90∘
Bond angles above are standard reference values; actual angles vary with lone-pair repulsion and the central atom/ligands (so use them as a starting point, not a memorised “exact” for every molecule).
CO2: linear molecule with dipoles that cancel.
NH3: trigonal pyramidal and polar due to lone pair.
2.2 Bond Polarity and Molecular Dipole
Even if bonds are polar, molecules can be non-polar when dipoles cancel (e.g. COX2). Use vector addition language: “the dipoles are equal and opposite along the linear axis.” For molecules such as NHX3, asymmetry yields a net dipole, linking to hydrogen bonding strength.
When writing IMF explanations, point to the trigonal pyramidal NHX3 sketch above: that shape gives a net dipole and the nitrogen lone pair enables hydrogen-bond interactions.
3 Intermolecular Forces (IMF)
IMF
Requirements
Relative strength
Example
London dispersion
Present in all molecules; induced dipoles.
Weakest; increases with molecular mass and surface area.
IX2
Permanent dipole-dipole
Polar molecules with permanent dipoles.
Medium.
CHX3Cl
Hydrogen bonding
H bonded to N, O, or F; lone pair on adjacent molecule.
Strongest IMF; drives high boiling points.
HX2O
Candidates should support IMF arguments with both electron distribution (polarity) and ability to form specific interactions (e.g. availability of lone pairs).
4 Worked Multi-Part Example
Question:
Consider PClX3 and PClX5.
Predict the shapes using VSEPR.
Explain why PClX5 exists whereas NClX5 does not.
Compare the melting points of PClX3 and PClX5
Solution outline:
PClX3: Steric number 4 (three bonding pairs + one lone pair) → trigonal pyramidal; bond angle slightly less than 109.5∘ due to lone-pair repulsion.
PClX5: Steric number 5 → trigonal bipyramidal.
Phosphorus forms a stable hypervalent PClX5 because its larger valence shell accommodates five bonding pairs with lower electron-pair repulsion (described by three-centre four-electron bonding); nitrogen is too small to stabilise five N−Cl bonds, so NClX5
PClX5 forms an ionic lattice in the solid state ([PClX4]X+[PClX6]X−
Each explanation links shape, orbital availability, and lattice structure-the exact reasoning markers want.
PCl3: trigonal pyramidal molecular species.
PCl5: framework used in shape and bonding comparisons.
5 Exam Tactics
Paper 1: Expect geometry MCQs where lone pairs must be counted correctly. Sketch quick Lewis structures to avoid errors.
Paper 2: Structured questions often request energy comparisons (e.g. lattice energies of MgO vs NaF). Reference ionic charge and ionic radius explicitly, stating relative charge density.
Paper 3: Long questions integrate bonding with properties. For example, deducing why silicon dioxide is a solid network while carbon dioxide remains gaseous at RTP-discuss covalent network vs discrete molecules.
Paper 4: Practical planning tasks may need justification of solvents based on polarity or explanation of observed conductivity based on ionic vs molecular structure.
6 Frequent Misconceptions
Confusing bond polarity with molecular polarity: Always consider shape.
Forgetting lone-pair repulsion: Bond angles decrease with each lone pair; quote specific angle adjustments.
Misstating metallic bonding explanations: Highlight delocalised electrons and cation lattice rather than vague “sea of electrons” statements.
Overcomplicating ‘expanded octet’ explanations: Keep to the syllabus framing (electron-pair repulsion + size/energetics) and write clearly rather than memorising contested orbital stories.
7 Quick Drill Set
Compare the melting points of MgClX2, NaCl, and AlX2OX3 using lattice energy arguments.
Sketch and label the shapes of SFX4, IFX5
Explain why graphite conducts electricity in the plane but diamond does not, referencing bonding and structure.
Worked solutions should mention the bonding model, structural arrangement, and the inter-particle forces in each case.
Common exam mistakes
Reporting molecular shape instead of electron-domain geometry: VSEPR requires you to state both - electron geometry first, then molecular shape after discounting lone pairs. Giving only the molecular shape misses marks when the question asks to "describe the shape".
Ignoring lone pairs when predicting bond angles: Bond angles in molecules such as NHX3 and HX2O are reduced below the tetrahedral ideal; failing to acknowledge lone-pair repulsion gives the wrong predicted angle.
Confusing bond polarity with molecular polarity: A molecule with polar bonds can still be non-polar if the bond dipoles cancel by symmetry (e.g. COX2, CClX4
Listing intermolecular forces without ranking strength: Questions that compare boiling points require you to identify the type of IMF and justify which is stronger; a bare list without comparative reasoning will not score explanation marks.
Overlooking London dispersion forces in polar molecules: All molecules experience London dispersion; omitting them when explaining properties of larger polar molecules gives an incomplete answer.
Mixing up sigma and pi bond counts: A C=C double bond contains one sigma and one pi bond; a triple bond contains one sigma and two pi bonds - reversing these is a common Paper 2 error.
Vague metallic bonding language: Phrases like "sea of electrons" alone are insufficient; state "a lattice of positive ions surrounded by delocalised electrons" and link this directly to the property asked about.
Frequently asked questions
Is VSEPR theory needed for every bonding question? VSEPR is the required model for predicting molecular geometry in H2 Chemistry. You are expected to apply the steric-number workflow for any molecule or ion with a central atom, and to link shape to polarity and intermolecular forces where relevant.
How do I decide which intermolecular force to discuss in a melting or boiling point comparison? First identify the types of IMF present in each substance. Then compare strength: hydrogen bonding > permanent dipole-dipole > London dispersion (though London forces can dominate for large non-polar molecules). Always justify which force is stronger rather than simply naming it.
What is the difference between lattice energy and bond energy? Lattice energy refers to the electrostatic energy of an ionic solid (enthalpy change when one mole of ionic compound forms from gaseous ions), while bond energy (bond dissociation enthalpy) refers to breaking a covalent bond in a gaseous molecule. They are assessed in different calculation types - Born-Haber cycles vs Hess' law bond-energy calculations.
Do I need to draw Lewis structures in the exam? Yes. Drawing a Lewis structure (showing all bonding pairs and lone pairs) is often the first step in VSEPR questions and is sometimes directly awarded marks. Always include lone pairs on the central atom and check the total valence electron count.
Struggling with Chemical Bonding? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
