1
Formal charge formula:
FC = V − L − S/2
V = valence electrons, L = lone pair electrons, S = shared (bonding) electrons
2
Ozone structure: O1=O2−O3
O1 (central): double bonded to O2, single bond to O3
Numbering in question: atom 1 = central, atoms 2 and 3 = terminal
Central O (atom 1): V=6, L=2 (1 lone pair), S=6 (double bond=4 + single bond=2... actually double bond counts as 4 shared). Wait — re-read: FC for central O in O3: V=6, L=2, S=8 (shares 4+4 for two bonds if double+single)... Let us recalculate carefully.
3
Ozone Lewis structure: O₂=O₁−O₃ (resonance hybrid)
One resonance structure: O₂=O₁−O₃
O₁ (central): 1 double bond to O₂, 1 single bond to O₃, 1 lone pair → V=6, L=2, S=6 → FC = 6−2−6/2 = 6−2−3 = +1
O₂ (double bonded terminal): V=6, L=4, S=4 → FC = 6−4−4/2 = 6−4−2 = 0
O₃ (single bonded terminal): V=6, L=6, S=2 → FC = 6−6−2/2 = 6−6−1 = −1
Ozone O₃: Formal charges → O(double bond terminal)=0, O(central)=+1, O(single bond terminal)=−1
Answer for atoms 2,1,3 → 0, +1, −1
1. Formal Charge Concept
FC = V − L − S/2. V=valence electrons (from group number). L=lone pair electrons on atom. S=shared electrons (bonding pairs × 2). FC tells how electrons are distributed in Lewis structure. Sum of FC in neutral molecule = 0. Sum of FC in ion = charge of ion. Preferred structure: minimise |FC|; negative FC on more electronegative atom.
2. Formal Charges in Ozone
Ozone has resonance: O=O−O ↔ O−O=O. In one resonance structure: central O has +1, double-bonded terminal O has 0, single-bonded terminal O has −1. In the other: reversed for the two terminals. The actual molecule is a resonance hybrid. Both resonance structures contribute equally — ozone has 1·5 bond order for each O-O bond.
3. Resonance in Ozone
Ozone structure: 3 O atoms, bent (117°). Not linear. The central O has 1 lone pair + 2 bonds. Due to resonance, both O-O bonds have equal length (128pm) — intermediate between O-O single (148pm) and O=O double (121pm). Resonance energy stabilises ozone. The actual structure cannot be represented by one Lewis structure.
4. VSEPR Geometry of Ozone
O₃: central O has 2 bond pairs + 1 lone pair = 3 electron domains → trigonal planar parent. With 1 lone pair: BENT (V-shape) geometry. Bond angle ~117° (less than 120° due to lone pair repulsion). Compare to SO₂: similar structure, similar bond angle (~119°).
5. Formal Charges vs Oxidation States
Formal charge: assumes equal sharing of bonding electrons (each atom gets S/2). Oxidation state: assumes complete transfer of electrons to more electronegative atom. For ozone: oxidation state of all O = 0 (pure element). But formal charges ≠ 0 for all atoms. These are different concepts. Formal charges help select the best Lewis structure; oxidation states are used in redox reactions.
6. Ozone — Properties and Uses
O₃ formed in upper atmosphere by UV splitting O₂: O₂ + UV → 2O•; O• + O₂ → O₃. Ozone layer (stratosphere, 15-35 km) absorbs harmful UV-B and UV-C. Ozone depletion by CFCs: CCl₂F₂ → Cl• (by UV); Cl• + O₃ → ClO• + O₂; ClO• + O• → Cl• + O₂ (chain reaction — one Cl• destroys ~100,000 O₃). O₃ is a powerful oxidising agent. Used in water purification, bleaching, disinfection.
7. Lewis Structures — Drawing Rules
1. Count total valence electrons. 2. Connect atoms with single bonds. 3. Place lone pairs to complete octets. 4. If octet not complete: form double/triple bonds. 5. Calculate formal charges. 6. Choose structure with: lowest formal charges, negative FC on electronegative atoms. Ozone: 18 valence electrons total (3×6). After connecting: O=O−O accounts for 4+2=6 bonding electrons, 12 remain for lone pairs.
8. Comparison: SO₂ vs O₃
SO₂: S=O−S structure (one form). S has formal charge +1 (using one double bond) or 0 (with expanded octet dπ-pπ bond). O₃ similar shape but all oxygen. Key difference: S can expand octet (d orbitals) → S=O bonds are shorter. O cannot expand octet → O₃ has equal 1·5-order bonds. Both are bent with similar bond angles (~117-119°). Both have resonance structures.
Frequently Asked Questions
1. How to calculate formal charge step by step? ⌄
For O₃, one resonance: O₂=O₁−O₃. Atom O₁ (central): V=6 (Group 16). Lone pairs on O₁=1 pair=2 electrons, so L=2. Bonds: 1 double bond (4 shared e⁻) + 1 single bond (2 shared e⁻) = 6 total shared. FC = 6 − 2 − 6/2 = 6 − 2 − 3 = +1. Atom O₂ (double bonded): V=6, L=4 (2 lone pairs), S=4. FC = 6−4−4/2 = 6−4−2 = 0. Atom O₃ (single bonded): V=6, L=6 (3 lone pairs), S=2. FC = 6−6−2/2 = 6−6−1 = −1.
2. Why is the central O atom +1 in ozone? ⌄
In the Lewis structure of ozone with one double bond and one single bond, the central O atom: loses 1 lone pair (has only 1 vs expected 2 for neutral O). The S/2 formula: shares 3 electrons effectively (half of 6 bonding electrons). But has only 2 non-bonding electrons (1 lone pair). FC = 6 − 2 − 3 = +1. The positive FC on central O reflects it sharing more electrons with neighbors than it 'owns'.
3. Why does the terminal O with single bond get −1? ⌄
Single-bonded terminal O in ozone: has 3 lone pairs (6 non-bonding electrons) and shares only 2 electrons in the single bond. FC = 6 − 6 − 2/2 = 6 − 6 − 1 = −1. It 'owns' more electrons (lone pairs) than neutral O would need. This negative FC reflects electron-rich character of the singly-bonded terminal oxygen — makes sense as it's more electronegative in this arrangement.
4. Is the actual ozone molecule better described by formal charges or resonance? ⌄
Neither single resonance structure is correct. The actual ozone molecule is a resonance hybrid of both: O₂=O₁−O₃ and O₂−O₁=O₃. Both contribute equally → both terminal O atoms have identical bond length (128 pm, intermediate). Both O-O bonds have order 1·5. The formal charges (+1 on central, and average of 0 and −1 = −0·5 on each terminal) in the resonance hybrid approximate the actual charge distribution.
5. What is the structure of SO₂ — similar to O₃? ⌄
SO₂: trigonal planar electron geometry, bent molecular geometry (~119°). One resonance structure: O=S−O (with lone pair on S). S(central): FC can be +1 (if S keeps 2 lone pairs) or 0 (if S uses dπ-pπ expanded octet). Since S is Period 3, it can expand octet. The dπ-pπ bonding makes SO₂ more stable than ozone. S-O bond length (143 pm) is intermediate between S-O single (160pm) and S=O double (143pm). Both O₃ and SO₂ are bent.
6. What is the formal charge on N in NO₃⁻? ⌄
NO₃⁻: One Lewis structure has N with 1 double bond to one O and 2 single bonds to two O atoms. N: V=5, L=0, S=8 (double bond 4 + 2 single bonds 2+2). FC(N) = 5−0−8/2 = 5−0−4 = +1. O with double bond: V=6, L=4, S=4. FC = 6−4−2 = 0. O with single bond: V=6, L=6, S=2. FC = 6−6−1 = −1 (each). Sum: +1 + 0 + (−1) + (−1) = −1 = charge of NO₃⁻ ✓
7. When should we prefer a structure with higher formal charges? ⌄
We should NOT prefer high formal charges. The best Lewis structure has: (1) formal charges as close to 0 as possible, (2) any negative FC on more electronegative atoms. For example, CO: with C triple bond to O and lone pair on each — C has FC=−1, O has FC=+1. Alternative with C=O (double bond): C=−2, O=+2 (worse). First option is better (smaller FC magnitudes). Exception: if octet expansion allows FC=0, prefer that for Period 3+ elements.
8. What is the ozone hole and how does it form? ⌄
Stratospheric ozone layer (15-50 km) shields Earth from UV radiation. CFCs (chlorofluorocarbons like Freon-12, CCl₂F₂) are stable in troposphere but broken down by UV in stratosphere: CCl₂F₂ → CF₂Cl• + Cl•. Chlorine radical chain reaction: Cl• + O₃ → ClO• + O₂; ClO• + O• → Cl• + O₂. Each Cl• can destroy ~10⁵ O₃ molecules. Antarctic ozone hole discovered 1985. Montreal Protocol (1987) banned CFCs. Replacements: HFCs (no Cl, but still greenhouse gases), HFOs. Ozone layer slowly recovering.