Topic 2 – Lewis Structures, Formal Charge & VSEPR

Valence Shell Fundamentals

• Valence electrons reside in the outermost shell; control bonding & chemical reactivity.

• Common student words from Mentimeter cloud: “outer shell”, “octet rule”, “shared”, “spdf”, “electronegativity”, “bond”, “chemical reaction”.

• Quick element recap (group 1 & 11 examples): Li, Na, K, Ag, Au all possess 1 e⁻ in their outer s-orbital.

Electronegativity (EN)

• Definition: numerical measure of an atom’s ability to attract shared electrons in a bond.

• Periodic trend → increases left-to-right across period, decreases top-to-bottom down a group.

• Self-rated comfort survey: 11 horrified, 20 worried, 30 comfortable, 9 very comfortable.

Lewis Structures – 5-Step Procedure

1. Count total valence electrons.

2. Assemble skeletal framework with single bonds (least electronegative atom central, never H).

3. Place three lone pairs (6 e⁻) on every outer atom (except H).

4. Assign remaining electrons to central atom.

5. Minimise formal charges (use multiple bonds/lone-pair shifts if needed).

Formal Charge Equation

{\text{Formal charge}}=\text{valence electrons}_{\text{free atom}}-\text{lone-pair electrons}-\frac{1}{2}(\text{bonding electrons})

Example 1: Phosgene Analogue, COCl$_2$

• Valence tally: C 4 + O 6 + 2Cl 14 = 24 e⁻.

• Framework uses 3 single bonds (6 e⁻) → 18 e⁻ remain.

• Distribute lone pairs: 3×(6 e⁻) on O, Cl, Cl = 18 e⁻ → 0 left.

• Initial formal charges: O = −1, C = +1, Cl = 0.

• Create C═O double bond; revised formal charges all 0.

Example 2: Boron Trichloride, BCl$_3$

• Valence: B 3 + 3Cl 21 = 24 e⁻.

• After 3 single B–Cl bonds: 18 e⁻ remain.

• Lone pairs on the 3 Cl atoms consume exactly 18 e⁻.

• Formal charges: B 0, each Cl 0; octet not achieved on B → electron-deficient but lowest FC.

Example 3: Azide Ion, N$_3^-$

• Target total e⁻ = 3×5 + 1 extra (−1 charge) = 16 e⁻.

• Multiple resonance forms; fully minimised set: terminal N −1, central N +1, satisfies octet via two N≡N/N═N patterns.

• Demonstrates need to recalc FC after every bond rearrangement.

Example Misdiagnosis: PO$_4^{3−}$

• Student error: assumed "5 bonds = 5 electron sets" (ignored lone pairs & charge) → predicted trigonal bipyramidal.

• Correct: 4 equivalent P–O bonds; tetrahedral electron set.

VSEPR (V​alence S​hell E​lectron P​air R​epulsion)

• Electron domains (bonding + lone pairs) orient to minimise repulsion.

• Repulsion hierarchy: \text{LP–LP} > \text{LP–BP} > \text{BP–BP}.

• LPs occupy more spatial volume → compress adjacent bond angles.

• Classic angle deviations (tetrahedral baseline 109.5^{\circ}):
  • NH$3$: 107^{\circ} (1 LP). • H$2$O: 104.5^{\circ} (2 LPs).

Geometry vs Shape

• Geometry = arrangement of all electron sets (bonding + lone pairs).

• Shape = geometry name after omitting the lone-pair positions.

• Mnemonic in slides: "GILP" (Geometry Includes Lone Pairs).

Master Table (also repeated throughout slides)

• 2 sets → Linear (shape always linear).

• 3 sets → Trigonal planar (shapes: trig. planar, bent).

• 4 sets → Tetrahedral (shapes: tetrahedral, trig. pyramidal, bent).

• 5 sets → Trigonal bipyramidal (shapes: trig. bipyramidal, seesaw, T-shaped, linear).

• 6 sets → Octahedral (shapes: octahedral, square pyramidal, square planar, T-shaped).

Worked VSEPR Examples

• NH$_3$: 4 electron sets → geometry tetrahedral, shape trigonal pyramidal.

• SO$_3^{2−}$: Central S has 4 electron sets (3 S–O bonds + 1 LP) → geometry tetrahedral, shape trigonal pyramidal.

• SF$_4$: 5 sets (4 BPs + 1 LP) → geometry trig. bipyramidal, shape seesaw.

• BrF$_5$: 6 sets (5 BPs + 1 LP) → geometry octahedral, shape square pyramidal.

• Heme cofactor: Fe centre changes from square planar (deoxygenated) to octahedral (oxygenated) upon O$_2$ binding.

Bond-Energy & Repulsion Illustration

• H–H interaction energy vs distance: curve minimum at bond length 74 pm.

• Bond energy 7.22 \times 10^{-19}\,\text{J} (≈436 kJ mol$^{-1}$).

Real-World & Biological Significance

• Shape determines polarity → e.g.
  • Linear "H–O–H" (180°) would not be water; bent shape is essential for water’s dipole & hydrogen bonding.

• ATP presented as polyphosphate example; multiple VSEPR centres govern its reactivity.

• Students asked to find VSEPR in favourite molecules (classroom activity).

Frequently Asked Questions (from polls)

• "How do I know if charges are minimised?" → all atoms closest to 0, negative charges on most electronegative atoms.

• "Which elements violate octet?" → B, Be often deficient; period 3+ atoms (P, S, Cl, Br, I) can expand octet.

• "Do we need to memorise bond angles?" → Learn ideal angles and typical deviations caused by 1 or 2 LPs.

• "Why can central N in azide have 4 bonds?" → uses dative/resonance; formal-charge bookkeeping more important than generic "N = 3 bonds" rule.

Equity & Professional Development

• Women-in-STEM session highlighted need for networking; encourages inclusive community building.

Summary Checklist for Exams

• Master 5-step Lewis procedure & formal-charge formula.

• Carry extra significant figures until final rounding.

• Memorise VSEPR table, repulsion hierarchy, and angle trends.

• Be able to move from Lewis structure → electron-set count → geometry → shape.

• Relate molecular shape to polarity, reactivity, and biological function.

• Practise with charged species & resonance (e.g., azide, phosphates).