Formal Charge - UNSOLVED PRACTICE SET
Chapter: Chemical Bonding and Molecular Structure | Topic: Formal Charge
FORMAL CHARGE - UNSOLVED PRACTICE SET
Topic: Formal Charge
Multiple Choice Questions
Q1. Formal charge on an atom in a Lewis structure is calculated as:
- Valence electrons − lone pair electrons
- Valence electrons − (lone pair electrons + ½ bonding electrons)
- Total electrons − valence electrons
- Bonding electrons − lone pair electrons
Q2. The most stable Lewis structure is the one in which:
- All atoms have zero formal charge
- Negative formal charges are on more electronegative atoms and positive formal charges on less electronegative atoms
- The central atom has the highest formal charge
- Formal charges are maximized
Q3. In the most stable Lewis structure of CO, the formal charges are:
- C = 0, O = 0
- C = +1, O = −1
- C = −1, O = +1
- C = +2, O = −2
Q4. The sum of formal charges in a neutral molecule is always:
- Equal to the number of atoms
- Zero
- Equal to the number of valence electrons
- Positive
Q5. In the Lewis structure of O₃, the central oxygen atom has a formal charge of:
- 0
- +1
- -1
- +2
Q6. Which of the following structures for NCO⁻ is most stable?
- [N≡C-O]⁻ with formal charges N = −1, C = 0, O = 0
- [N=C=O]⁻ with formal charges N = 0, C = 0, O = −1
- [⁻N-C≡O] with formal charges N = −2, C = +1, O = 0
- All are equally stable
Short Answer Questions
Q7. Define formal charge. Write the formula for calculating formal charge on an atom in a molecule.
Q8. Calculate the formal charge on each atom in the Lewis structure of CO₂.
Q9. Why is the structure with the smallest formal charges (preferably zero) on all atoms considered the most stable Lewis structure?
Q10. In the Lewis structure of HNC (hydrogen isocyanide), hydrogen is bonded to nitrogen. Calculate the formal charges and compare with HCN. Which is more stable?
Q11. Your father is reviewing the family budget. He says the most stable financial situation is when everyone's share balances out to zero — no one owes anyone. How is this exactly analogous to the concept of formal charge in molecules, where the most stable structure has formal charges closest to zero on all atoms?
Q12. Calculate the formal charges on all atoms in the sulphate ion (SO₄²⁻) for the structure with two S=O double bonds and two S-O single bonds.
Long Answer Questions
Q13. Explain the concept of formal charge and its significance in determining the most stable Lewis structure. Describe the rules for selecting the best resonance structure based on formal charges. Calculate the formal charges for all atoms in the following structures and identify the most stable:
(a) OCN⁻ (cyanate ion)
(b) NCO⁻ (isocyanate ion)
(c) CNO⁻ (fulminate ion
Q14. Discuss how formal charge helps in understanding the actual electronic distribution in molecules. For the nitrous oxide molecule (N₂O), draw three possible Lewis structures and calculate formal charges for each. Which structure is the major contributor to the resonance hybrid? Why? What does this tell us about the N-N and N-O bond orders in N₂O?
Q15. During a chemistry competition, you are given the following challenge:
(a) Draw three possible Lewis structures for the cyanate ion (OCN⁻) and calculate formal charges for each atom in each structure.
(b) Apply the rules: (i) formal charges closest to zero are preferred, (ii) negative formal charge on more electronegative atom is preferred. Which structure wins?
(c) For the thiocyanate ion (SCN⁻), the structure [S-C≡N]⁻ is more stable than [⁻S-C≡N]. Explain using formal charges and electronegativity.
(d) Why does CO have a structure with a triple bond and formal charges (C⁻, O⁺) rather than a double bond with zero formal charges? What does this tell us about the limitations of formal charge analysis?
Numerical / Application-Based Problems
Q16. Calculate the formal charge on each atom in the following Lewis structures:
(a) Ozone (O₃): central O bonded to two terminal O atoms, with one O=O double bond and one O-O single bond
(b) Nitrate ion (NO₃⁻): central N with one N=O double bond and two N-O single bonds
(c) Carbon monoxide (CO): C≡O with one lone pair on C and one lone pair on O
For each, verify that the sum of formal charges equals the overall charge of the species.
Q17. For the phosphate ion (PO₄³⁻), consider two Lewis structures:
Structure A: One P=O double bond and three P-O single bonds
Structure B: All P-O single bonds with expanded octet on phosphorus
(a) Calculate formal charges for all atoms in Structure A.
(b) Explain why Structure B (with zero formal charges on all oxygens) is not typically drawn, even though it minimizes formal charges.
(c) The actual structure is a resonance hybrid. What is the average P-O bond order in PO₄³⁻?
(d) Compare the P-O bond lengths in PO₄³⁻ with the P=O bond length in P₄O₁₀. Which is shorter and why?
Q18. Consider the following molecules/ions and their possible Lewis structures:
(a) SO₂: Draw two resonance structures and calculate formal charges. Which structure contributes more to the hybrid?
(b) ClO₃⁻: Draw the Lewis structure and calculate formal charges. Why does chlorine have an expanded octet?
(c) For H₂SO₄, draw the structure and calculate formal charges on all atoms. Why does sulphur have a +2 formal charge in the most common representation?
(d) Compare the S-O bond lengths in H₂SO₄ and SO₄²⁻. Explain any differences using formal charge and resonance concepts.