Lewis structures — count valence electrons, draw bonds and lone pairs. The blueprint of every molecule.
Act II
What shape do they build?
VSEPR theory — electron domains form a parent geometry; lone pairs compress it into predictable shapes and angles.
Act III
Is the molecule polar?
Electronegativity + symmetry — bond dipoles may cancel or combine. This decides how a molecule behaves.
◈The exam logic: draw the Lewis structure → determine the shape → check bond dipoles & symmetry → decide polar or non-polar. Every question is this chain.
01
03 / 10
Act I · Lewis structures
Drawing Lewis structures — the 5-step method
Count total valence electrons. Add one per negative charge, subtract one per positive charge.
Place the least electronegative atom central (H is never central). Connect with single bonds.
Complete outer octets of the surrounding atoms with lone pairs.
Leftover electrons go on the central atom — as lone pairs.
If the central atom lacks an octet, make multiple bonds — turn lone pairs into double/triple bonds.
Watch out
Common traps
✕ Forgetting to adjust for charge — NH₄⁺ has 8 valence e⁻ (5 + 4 − 1), not 9.
✕ Hydrogen only needs 2 electrons — never an octet.
✕ Boron & beryllium can be electron-deficient (BF₃ is fine with 6 e⁻).
✎Exam habit: show lone pairs as dots and bonding pairs as lines — and label your final answer with the total electron count.
02
04 / 10
Act I · Worked examples
Lewis structures in practice.
H₂O — 8 e⁻ total
2 bonding pairs · 2 lone pairs
CO₂ — 16 e⁻ total
no lone pairs on C
NH₄⁺ — 8 e⁻ total
5 + 4 − 1 = 8 valence e⁻
⚛Notice: the lone pairs and multiple bonds you draw here directly determine the shape on the next slides.
03
05 / 10
Act II · Shape of molecules
Every shape starts from its parent geometry
VSEPR: the electron domains around the central atom (bonding pairs + lone pairs) repel and spread as far apart as possible. The arrangement they form is the parent (electron) geometry — it sets the starting bond angle. Lone pairs then compress the observed angle below the parent value.
Electron domains
Parent geometry
Parent angle
2
Linear
180°
3
Trigonal planar
120°
4
Tetrahedral
109.5°
Repulsion strength: LP–LP > LP–BP > BP–BP — lone pairs sit closer to the nucleus and squeeze bonding pairs together.
Same parent · shrinking angle
All three have 4 domains → tetrahedral parent of 109.5°
CH₄ · 4 BP, 0 LP109.5°
NH₃ · 3 BP, 1 LP≈ 107°
H₂O · 2 BP, 2 LP≈ 104.5°
Rule of thumb: each lone pair costs about 2.5° — but in the exam, always explain it as "less than 109.5° because lone pairs repel more strongly".
◎Parent first, shape second: the electron geometry is tetrahedral in all three — the molecular shape is named only by where the atoms end up sitting.
04
06 / 10
Act II · Learn these six
Shapes & bond angles — explained from the parent angle
Molecule
Parent geometry (ideal angle)
Molecular shape
Bond angle
Why the angle is what it is
CH₄
Tetrahedral · 109.5°
Tetrahedral
109.5°
4 bonding pairs — repulsion is equal, no lone pairs, so it stays at the parent angle
NH₃
Tetrahedral · 109.5°
Trigonal pyramidal
≈ 107°
1 lone pair repels the bonding pairs more strongly, compressing 109.5° → ≈107°
H₂O
Tetrahedral · 109.5°
Bent (V-shaped)
≈ 104.5°
2 lone pairs repel the bonding pairs even more, compressing 109.5° → ≈104.5°
BF₃
Trigonal planar · 120°
Trigonal planar
120°
Only 3 bonding pairs, no lone pairs — stays at the parent 120°
CO₂
Linear · 180°
Linear
180°
2 double bonds act as 2 domains — stays at the parent 180° (and is symmetric → non-polar)
PCl₅ (ext.)
Trig. bipyramidal · 90°/120°
Trigonal bipyramidal
90° / 120°
5 bonding pairs — extension beyond the 4-domain family
✍Model NCEA answer: "The 4 electron pairs around the central atom form a tetrahedral arrangement based on 109.5°. The lone pair(s) repel more strongly than bonding pairs and are held closer to the central atom, pushing the bonding pairs closer together — so the bond angle is less than 109.5° (≈107° for NH₃, ≈104.5° for H₂O)."
05
07 / 10
Act III · Polarity part 1
Polar bonds — the tug-of-war for electrons
Atoms share electrons unequally when their electronegativity differs. The more electronegative atom pulls the shared pair closer — creating a dipole.
The electron cloud shifts toward chlorine. The bond is polar — and a diatomic molecule with a polar bond is always a polar molecule.
⇌Say it properly in the exam: "Cl is more electronegative than H, so the bonding electrons are attracted more strongly to Cl, giving Cl a partial negative charge (δ⁻) and H a partial positive charge (δ⁺)."
06
08 / 10
Act III · Polarity part 2
Polar vs non-polar molecules.
Polar molecule
Asymmetric — dipoles don't cancel
Bent shape + polar O–H bonds → the dipole arrows cannot cancel. Polar molecule.
net dipole ≠ 0
vs
Non-polar molecule
Symmetric — dipoles cancel
Each C=O bond is polar, but the linear symmetry makes the dipoles cancel. Non-polar molecule.