Resonance, VSEPR Shapes & Molecular Polarity – Comprehensive Study Notes

Lewis theory assumes electrons are localised: either
- Shared between two specific atoms (bonding pairs)
- Un-shared on one atom (lone pairs)
Certain ions/molecules show experimental data that contradict the localised picture
- Example numbers written on board (peak bond lengths in pm): 946 pm → 418 pm → 163 pm: trend shows bond length decreases with higher bond order
Classic failure: \mathrm{NO_2^-}
- Straightforward Lewis drawing gives one N–O double bond (short) & one N–O single bond (long)
- Experimental fact: both N–O bonds are identical in length; no long/short pair
- Conclusion: electrons are delocalised → need Resonance

Represent delocalised bonding by sketching all equivalent Lewis variants connected by a double-headed arrow ⇌
Rules
- Same atom skeleton; only electron placement differs
- True structure is the average (hybrid) of all resonance forms
For \mathrm{NO_2^-}
- Two contributing forms
- Form 1: left O double-bonded
- Form 2: right O double-bonded
- Bond order per N–O = \dfrac{1+2}{2}=1.5
Resonance vs isomers
- Isomers = different compounds (atoms re-arranged)
- Resonance = same compound, electrons shuffled
Other resonance examples mentioned
- Benzene \mathrm{C6H6}: alternating three C=C → hybrid has no true double bonds, every C–C = bond order 1.5; confirmed because bromine test shows no addition reaction
- Polyatomic chlorite/chlorate/perchlorate: electron pairs on Cl form coordinate covalent bonds (all valid resonance)

Count total valence e⁻ (add for negative charge, subtract for positive)
Choose central atom (usually lowest EN, never H)
Connect atoms with single bonds; distribute remaining e⁻ to complete octets on outer atoms, then central
If e⁻ run short, introduce π-bonds or formal charges
For ions/molecules with multiple equivalent positions, draw all resonance forms; test by shifting electron pairs, NOT atoms

Shared pair comes exclusively from one atom (e.g.
- ClO⁻, ClO₂⁻, ClO₃⁻, ClO₄⁻ → all extra pairs donated by Cl)
Still counted as a single bond in VSEPR

Concept: Electron groups (lone pairs & all bonds – single, double, triple each count as 1 group) repel to maximise separation.

Electron Groups	Basic Geometry	Ideal Angles	Board example
2	Linear	180^\circ	\mathrm{CO_2}
3	Trigonal planar	120^\circ	\mathrm{CO_3^{2-}}
4	Tetrahedral	109.5^\circ	\mathrm{CH_4}
+1 lone pair (3 + 1)	Trigonal pyramidal	\approx107^\circ	\mathrm{NH_3}
+2 lone pairs (2 + 2)	Bent / V-shaped	\approx104.5^\circ	\mathrm{H_2O}
(Simple lecture stopped at 4 groups; higher geometries — trigonal bipyramidal, octahedral — exist but not yet covered.)

\mathrm{CO_2}
- Central C attached to 2 O, 0 lone pairs → 2 e-groups → linear
\mathrm{CO_3^{2-}}
- Central C + 3 O → 3 groups → trigonal planar (all atoms coplanar)
\mathrm{CH_4}
- 4 bonds, 0 LP → tetrahedral; three H in plane, one out/dash technique
\mathrm{NH_3}
- 3 bonds + 1 LP → trigonal pyramidal; LP occupies one vertex of tetrahedron → angle compressed to ~107^\circ

Electronegativity (EN) = ability to attract shared e⁻; scale: F highest (~4.0)
ΔEN = |EN₁ – EN₂| determines bond type
- \Delta\text{EN} \le 0.4 → non-polar covalent
- 0.5 \le \Delta\text{EN} \le 1.9 → polar covalent (partial charges, dipole arrow toward more EN atom)
- \Delta\text{EN} \ge 2.0 → ionic (no sharing; electrostatic attraction)
Dipole moment vector: arrow from δ⁺ toward δ⁻ ; cross on positive end

Process

Diatomics (2 atoms): polarity = bond polarity directly
- \mathrm{NO} (ΔEN≈0.5) → polar
- \mathrm{O2}, \mathrm{Cl2} → non-polar
Water \mathrm{H_2O}
- Bent shape, two equal dipoles at ≈104.5° → do not cancel ⇒ net dipole → polar (explains high solubility properties)
Carbon dioxide \mathrm{CO_2}
- Linear; two identical C→O dipoles 180° apart cancel exactly ⇒ non-polar (despite polar bonds)

Ion	Total valence e⁻	LP on Cl (after octets)	Note
\mathrm{ClO^-}	14	3	1 coordinate bond
\mathrm{ClO_2^-}	20	2	resonance among O positions
\mathrm{ClO_3^-}	26	1	trigonal pyramidal (sp³)
\mathrm{ClO_4^-}	32	0	tetrahedral, all Cl–O equivalent
(All extra bonds form from lone pairs on Cl → "coordinate covalent")

Always finish Lewis structures before VSEPR or polarity steps
For multi-centre molecules (e.g. \mathrm{C2H2}) treat each central atom separately for geometry
Diatomic molecules: list as linear, count 1 bond, then evaluate polarity directly
Worksheet columns (at minimum)
1. Total e-groups on central atom
2. Molecular shape name
3. Ideal bond angle(s)
4. Polarity (last; depends on dipole cancellation)
Use molecular model kits to visualise 3-D; helpful for wedge/dash drawings
Recommended to draw in pencil; avoid accidental 90° angles unless genuine

Progressive bond length example: 946 pm → 418 pm → 163 pm (illustrates inverse relation to bond order)
Ideal tetrahedral angle: 109.5^{\circ}; trigonal pyramidal (NH₃) compressed to \approx107^{\circ}; H₂O bent ≈104.5^{\circ}
Bond order calculation via resonance average: e.g. \mathrm{NO_2^-} yields 1.5
Bromine test for unsaturation: orange Br₂ decolourises with true C=C; benzene shows no reaction (supports resonance model)

Localised Lewis structures sometimes fail → resonance needed
Resonance: same atomic framework, electron redistribution; actual bond order = average
Coordinate covalent = both electrons supplied by one donor atom but treated as normal bond in VSEPR
VSEPR uses total electron groups to predict 3-D shapes; lone pairs modify ideal angles
Bond polarity from ΔEN; molecular polarity from vector sum of bond dipoles
Geometry, resonance & polarity underpin physical properties (solubility, reactivity, spectroscopy)