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IB DP Chemistry – S2.2.14 Formal Charge and Lewis Formulas- Study Notes

IB DP Chemistry - S2.2.14 Formal Charge and Lewis Formulas- Study Notes - New Syllabus - 2026, 2027 & 2028

IB DP Chemistry – S2.2.14 Formal Charge and Lewis Formulas- Study Notes – New Syllabus

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Structure 2.2.14 — Formal Charge and Its Use in Evaluating Lewis Structures

Structure 2.2.14 — Formal Charge and Its Use in Evaluating Lewis Structures

Formal Charge:

The formal charge (FC) of an atom in a molecule or ion is a hypothetical charge assigned to that atom, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of electronegativity.
It is a tool used to identify the most stable or preferred Lewis structure among possible alternatives.

Formula for Formal Charge:

    

\( \text{Formal Charge (FC)} = \text{Valence electrons} – \text{Non-bonding electrons} – \frac{1}{2} \times \text{Bonding electrons} \)

  • Valence electrons: from the periodic table (group number)
  • Non-bonding electrons: lone pair electrons on the atom
  • Bonding electrons: shared electrons (in bonds) attached to that atom

Criteria for the Best Lewis Structure:

  • Lowest number of atoms with non-zero formal charges
  • Negative formal charges on more electronegative atoms
  • Overall formal charge adds up to the molecule/ion’s total charge

Example: Formal Charge in Ozone \( \text{O}_3 \)

There are two major resonance structures of ozone:

  1. Double bond on the left: \( \text{O} = \text{O} – \text{O} \)
  2. Double bond on the right: \( \text{O} – \text{O} = \text{O} \)

So the formal charges are: +1, 0, -1 → Total charge = 0 

The true structure is a resonance hybrid of both forms.

Tips:

  • Use formal charges to compare possible Lewis structures
  • Avoid placing like charges adjacent to each other
  • Zero or small formal charges indicate a more likely structure
  • Never place a positive charge on the most electronegative atom unless justified

Example

Which Lewis structure is better for nitrate ion \( \text{NO}_3^- \)?

▶️ Answer/Explanation

There are three possible resonance structures, each with a different position for the double bond.

  • N atom: 5 valence electrons
  • Bonding: 1 double bond (4e⁻), 2 single bonds (4e⁻) → total 8 bonding e⁻ → 4 bonding pairs
  • No lone pairs on N → Formal charge = \( 5 – 0 – \frac{8}{2} = +1 \)

For one O with a double bond:

  • 6 valence e⁻, 4 non-bonding, 4 bonding → FC = \( 6 – 4 – \frac{4}{2} = 0 \)

For each O with single bonds:

  • 6 valence e⁻, 6 non-bonding, 2 bonding → FC = \( 6 – 6 – \frac{2}{2} = -1 \)

So charges are: N = +1, 1 double-bond O = 0, 2 single-bond O = -1 each
Total = -1 (✔ matches charge of ion)

Conclusion: All 3 resonance structures are equivalent. True structure is the average (delocalized).

Example

Choose preferred Lewis structure for \( \text{CO}_2 \)

▶️ Answer/Explanation

  1. \( \text{O} = \text{C} = \text{O} \) (double bonds)
  2. \( \text{O} – \text{C} \equiv \text{O} \) (one single, one triple)

Structure 1: Double bonds

  • C: 4 valence e⁻, 0 lone pairs, 4 bonding pairs → FC = \( 4 – 0 – 4 = 0 \)
  • O: 6 valence e⁻, 4 non-bonding e⁻, 4 bonding e⁻ → FC = \( 6 – 4 – 2 = 0 \)

Structure 2: Single and triple bond

  • C: 4 valence e⁻, 0 lone pairs, 5 bonding pairs → FC = \( 4 – 0 – 5 = -1 \)
  • O atoms: one FC = +1, other = 0

Conclusion: First structure is better → formal charges are all zero.

Example

Formal charges in sulfur dioxide \( \text{SO}_2 \)

Multiple valid Lewis structures exist. Which is preferred?

▶️ Answer/Explanation

Structure: 

  • S: 6 valence e⁻, 2 lone e⁻, 6 bonding e⁻ → FC = \( 6 – 2 – 3 = +1 \)
  • Double-bonded O: FC = 0
  • Single-bonded O: 6 non-bonding e⁻ + 2 bonding → FC = \( 6 – 6 – 1 = -1 \)

Charges: S = +1, one O = 0, one O = -1 → total = 0 ✔

Conclusion: Resonance between two such structures with delocalization of double bond.

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