Lewis Structures Practice Worksheet - Free Printable

Since I can't view or access uploaded images directly, I can’t see the specific task or image you’re referring to. However, based on your description, it seems like you’ve shared a Lewis Dot Structure Worksheet for AP Chemistry that includes a list of compounds and ions (e.g., SO₃, CO₃²⁻, NH₃, etc.), and the goal is to:

1. Draw Lewis structures for each compound/ion.
2. Calculate formal charges for atoms in the structure.
3. Observe octet rule compliance and exceptions.

I’ll walk you through how to solve this type of problem systematically using general principles, and then provide examples for several compounds from your list. You can apply the same method to the rest.

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🔧 Step-by-Step Guide: Drawing Lewis Structures & Calculating Formal Charges

#### ✔ Step 1: Count total valence electrons
Use the periodic table:
- Group 1A → 1 valence e⁻
- Group 2A → 2
- Group 3A → 3
- Group 4A → 4
- Group 5A → 5
- Group 6A → 6
- Group 7A → 7
- Group 8A → 8
- Add or subtract electrons for charges (anions: +e⁻, cations: –e⁻)

#### ✔ Step 2: Identify central atom
- Usually the least electronegative atom (except H, which goes on outside).
- For polyatomic ions, often the central atom is the one with lower electronegativity (e.g., S in SO₄²⁻, C in CO₃²⁻).

#### ✔ Step 3: Draw single bonds first
- Each bond = 2 electrons
- Connect outer atoms to central atom

#### ✔ Step 4: Distribute remaining electrons
- Fill octets (8 electrons) on outer atoms first (H gets 2)
- Then place lone pairs on central atom

#### ✔ Step 5: Check octet rule
- If central atom lacks octet → form double/triple bonds
- Use formal charge to evaluate best structure

#### ✔ Step 6: Calculate formal charge (FC)
\[
\text{FC} = \text{Valence e⁻} - (\text{Lone pair e⁻} + \frac{1}{2} \times \text{Bonding e⁻})
\]

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Now let’s go through several examples from your list:

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🌟 Example 1: SO₃ (Sulfur trioxide)

Step 1: Valence electrons
- S: 6
- O × 3: 6 × 3 = 18
- Total = 6 + 18 = 24 e⁻

Step 2: Central atom → S (less electronegative than O)

Step 3: Single bonds
- S–O × 3 → 3 × 2 = 6 e⁻ used

Step 4: Remaining electrons
- 24 – 6 = 18 e⁻ → place on O atoms (each O needs 6 more e⁻ to complete octet)
- 3 O atoms × 6 e⁻ = 18 → all used

But now S has only 6 e⁻ (3 bonds), so double bonds needed

Try resonance:
- One double bond between S and one O, two single bonds
- But better: three resonance structures where S forms double bonds with each O (expanded octet allowed for S)

Final Lewis structure: S with three double bonds to O atoms (no lone pairs on S). Each O has two lone pairs.

Formal charge on S:
- Valence = 6
- Lone pairs = 0 → 0 e⁻
- Bonding e⁻ = 6 bonds × 2 = 12 → half = 6
- FC = 6 – (0 + 6) = 0

Formal charge on O (double-bonded):
- Valence = 6
- Lone pairs = 4 e⁻
- Bonding e⁻ = 4 → half = 2
- FC = 6 – (4 + 2) = 0

✔ All formal charges = 0 → good!

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🌟 Example 2: CO₃²⁻ (Carbonate ion)

Step 1: Valence electrons
- C: 4
- O × 3: 6 × 3 = 18
- Charge: +2 e⁻ → total = 4 + 18 + 2 = 24 e⁻

Step 2: Central atom → C

Step 3: Bonds
- C–O × 3 → 6 e⁻ used

Step 4: Remaining electrons
- 24 – 6 = 18 e⁻ → place on O atoms (each O gets 6 e⁻ → 3 lone pairs per O)

But now C has only 6 e⁻ → need double bond

Resonance structure:
- One C=O double bond, two C–O⁻ single bonds (with negative charge on O)
- Three resonance structures (double bond rotates among O atoms)

Formal charges:
- C: Valence = 4; Lone pairs = 0; Bonding e⁻ = 8 (4 bonds) → FC = 4 – (0 + 4) = 0
- Double-bonded O: Valence = 6; Lone pairs = 4 e⁻; Bonding e⁻ = 4 → FC = 6 – (4 + 2) = 0
- Single-bonded O⁻: Valence = 6; Lone pairs = 6 e⁻; Bonding e⁻ = 2 → FC = 6 – (6 + 1) = –1

Total charge = 0 + 0 + (–1) × 2 = –2 → matches ion charge ✔

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🌟 Example 3: NH₃ (Ammonia)

Valence electrons:
- N: 5
- H × 3: 1 × 3 = 3
- Total = 8 e⁻

Structure:
- N central, bonded to 3 H atoms → 3 bonds = 6 e⁻
- Remaining 2 e⁻ → lone pair on N

Formal charge on N:
- Valence = 5
- Lone pairs = 2 e⁻
- Bonding e⁻ = 6 → half = 3
- FC = 5 – (2 + 3) = 0

Each H: FC = 1 – (0 + 1) = 0

✔ Octet complete on N (8 e⁻), H has duet.

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🌟 Example 4: CH₄ (Methane)

- C: 4, H × 4: 4 → total 8 e⁻
- C bonded to 4 H → 4 single bonds
- No lone pairs
- FC on C: 4 – (0 + 4) = 0
- FC on H: 1 – (0 + 1) = 0

✔ All good.

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🌟 Example 5: H₂O

- O: 6, H × 2: 2 → total 8 e⁻
- O bonded to 2 H → 4 e⁻ used
- Remaining 4 e⁻ → 2 lone pairs on O
- FC on O: 6 – (4 + 2) = 0
- H: 0

✔ Octet on O, duet on H.

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🌟 Example 6: NH₄⁺ (Ammonium ion)

- N: 5, H × 4: 4 → total 9 e⁻, but +1 charge → subtract 1 → 8 e⁻
- N bonded to 4 H → 4 bonds = 8 e⁻
- No lone pairs on N
- FC on N: 5 – (0 + 4) = +1
- Each H: 0
- Total charge = +1 → correct

⚠️ Exception: N has 8 e⁻ (octet), but formal charge +1.

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🌟 Example 7: NO₃⁻ (Nitrate ion)

- N: 5, O × 3: 18, +1 e⁻ (from –1 charge) → total = 24 e⁻
- Central atom: N
- Resonance: One N=O double bond, two N–O⁻ single bonds
- Three resonance structures

Formal charges:
- N: 5 – (0 + 4) = +1
- Double-bonded O: 6 – (4 + 2) = 0
- Single-bonded O⁻: 6 – (6 + 1) = –1
- Total charge: +1 + 0 + (–1)×2 = –1 → correct

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🌟 Example 8: SO₄²⁻ (Sulfate ion)

- S: 6, O × 4: 24, +2 e⁻ → total 32 e⁻
- Central S, four O atoms
- Two S=O double bonds, two S–O⁻ single bonds (resonance)
- S has expanded octet (12 e⁻)
- FC on S: 6 – (0 + 6) = 0
- Double-bonded O: FC = 0
- Single-bonded O: FC = –1
- Total charge = –2 → correct

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🔁 General Tips for Resonance and Exceptions

- Resonance: When multiple valid Lewis structures exist, draw them with double-headed arrows.
- Octet exceptions:
- Expanded octets: Elements in period 3+ (P, S, Cl, Xe) can have >8 e⁻
- Incomplete octets: BF₃, BeCl₂
- Odd electron molecules: NO, NO₂

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✔ Summary Table for Quick Reference

| Compound | Total e⁻ | Central Atom | Formal Charges | Notes |
|--------|---------|--------------|----------------|-------|
| SO₃ | 24 | S | S=0, O=0 | Resonance, expanded octet |
| CO₃²⁻ | 24 | C | C=0, O=0 or –1 | Resonance |
| NH₃ | 8 | N | N=0 | Octet complete |
| CH₄ | 8 | C | All 0 | Tetrahedral |
| H₂O | 8 | O | O=0 | Bent shape |
| NH₄⁺ | 8 | N | N=+1 | Positive ion |
| NO₃⁻ | 24 | N | N=+1, O=0 or –1 | Resonance |
| SO₄²⁻ | 32 | S | S=0, O=0 or –1 | Expanded octet |

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📝 How to Complete Your Worksheet

For each compound:
1. Count valence electrons
2. Draw skeleton (central atom)
3. Add bonds and lone pairs
4. Check octets
5. Adjust with double/triple bonds if needed
6. Calculate formal charges
7. Draw resonance if applicable

Let me know if you'd like detailed Lewis structures drawn out for any specific compound (e.g., H₂O₂, C₂H₄, PO₄³⁻, etc.) — I can describe them clearly.

Would you like me to go through all 33 compounds step by step? Or focus on a few tricky ones like H₂O₂, C₂H₄, CH₃OH, or PO₄³⁻?