Worksheet for practicing Lewis structures of various molecules and ions, with instructions to include resonance structures and determine preferred structures using formal charges.

I'll help you solve these Lewis structure problems step by step. Since I can't draw images, I'll describe each Lewis structure in detail and explain the reasoning.

## General Steps for Drawing Lewis Structures:
1. Count total valence electrons
2. Place atoms (least electronegative in center)
3. Connect with single bonds
4. Distribute remaining electrons to satisfy octets
5. Check formal charges and consider resonance

Let me go through each one:

1. CH₄ (Methane)

- C has 4 valence e⁻, each H has 1 → Total: 8 e⁻
- C in center, bonded to 4 H atoms
- All atoms have complete octets (H has duet)
- Structure: C with 4 single bonds to H

2. BF₃ (Boron trifluoride)

- B has 3 valence e⁻, each F has 7 → Total: 24 e⁻
- B in center, bonded to 3 F atoms
- B has only 6 electrons (incomplete octet), but this is stable due to boron's ability to be electron-deficient
- Structure: B with 3 single bonds to F, each F has 3 lone pairs

3. HCN (Hydrogen cyanide)

- H: 1, C: 4, N: 5 → Total: 10 e⁻
- H-C≡N (triple bond between C and N)
- C has 4 bonds (1 to H, 3 to N), N has 3 bonds + 1 lone pair
- Structure: H-C≡N with lone pair on N

4. NH₃ (Ammonia)

- N: 5, each H: 1 → Total: 8 e⁻
- N in center, bonded to 3 H atoms
- N has 1 lone pair
- Structure: N with 3 single bonds to H and 1 lone pair

5. H₂O (Water)

- O: 6, each H: 1 → Total: 8 e⁻
- O in center, bonded to 2 H atoms
- O has 2 lone pairs
- Structure: O with 2 single bonds to H and 2 lone pairs

6. ClO₃⁻ (Chlorate ion)

- Cl: 7, each O: 6, +1 for negative charge → Total: 26 e⁻
- Cl in center, bonded to 3 O atoms
- Resonance structures exist with double bonds
- Most stable: One Cl=O double bond, two Cl-O⁻ single bonds with formal charges
- Formal charges: Cl = +2, O (double bond) = 0, O (single bond) = -1 each

7. H₃O⁺ (Hydronium ion)

- O: 6, each H: 1, -1 for positive charge → Total: 8 e⁻
- O in center, bonded to 3 H atoms
- O has 1 lone pair
- Structure: O with 3 single bonds to H and 1 lone pair

8. NH₂⁻ (Ammide ion)

- N: 5, each H: 1, +1 for negative charge → Total: 8 e⁻
- N in center, bonded to 2 H atoms
- N has 2 lone pairs
- Structure: N with 2 single bonds to H and 2 lone pairs

9. NH₄⁺ (Ammonium ion)

- N: 5, each H: 1, -1 for positive charge → Total: 8 e⁻
- N in center, bonded to 4 H atoms
- No lone pairs on N
- Structure: N with 4 single bonds to H

10. C₂H₄ (Ethylene)

- Each C: 4, each H: 1 → Total: 12 e⁻
- Two C atoms double bonded to each other
- Each C bonded to 2 H atoms
- Structure: H₂C=CH₂

11. C₂H₂ (Acetylene)

- Each C: 4, each H: 1 → Total: 10 e⁻
- Two C atoms triple bonded to each other
- Each C bonded to 1 H atom
- Structure: HC≡CH

12. N₃⁻ (Azide ion)

- Each N: 5, +1 for negative charge → Total: 16 e⁻
- Resonance structures: [N=N⁺=N]⁻ ↔ [N⁻=N⁺=N]
- Most stable: Central N has +1 formal charge, terminal N's have -1 and 0
- Structure: Resonance hybrid with N=N=N, one double bond, one triple bond

13. NOF (Nitrosyl fluoride)

- N: 5, O: 6, F: 7 → Total: 18 e⁻
- N in center, bonded to O and F
- Likely N=O double bond, N-F single bond
- N has formal charge of +1, O has -1, F has 0
- Structure: F-N=O with lone pairs

14. ClF₂⁺ (Chlorine difluoride cation)

- Cl: 7, each F: 7, -1 for positive charge → Total: 20 e⁻
- Cl in center, bonded to 2 F atoms
- Cl has expanded octet (10 electrons)
- Structure: Cl with 2 single bonds to F, 2 lone pairs on Cl

15. FNO₂ (Nitryl fluoride)

- F: 7, N: 5, each O: 6 → Total: 24 e⁻
- N in center, bonded to F and 2 O atoms
- Likely N=O double bonds, N-F single bond
- Structure: F-N(=O)=O with lone pairs

16. XeF₅⁺ (Xenon pentafluoride cation)

- Xe: 8, each F: 7, -1 for positive charge → Total: 40 e⁻
- Xe in center, bonded to 5 F atoms
- Xe has expanded octet
- Structure: Xe with 5 single bonds to F, 1 lone pair

17. BrF₅ (Bromine pentafluoride)

- Br: 7, each F: 7 → Total: 42 e⁻
- Br in center, bonded to 5 F atoms
- Br has expanded octet
- Structure: Br with 5 single bonds to F, 1 lone pair

18. IF₄⁻ (Iodine tetrafluoride anion)

- I: 7, each F: 7, +1 for negative charge → Total: 36 e⁻
- I in center, bonded to 4 F atoms
- I has expanded octet
- Structure: I with 4 single bonds to F, 2 lone pairs

19. SF₆ (Sulfur hexafluoride)

- S: 6, each F: 7 → Total: 48 e⁻
- S in center, bonded to 6 F atoms
- S has expanded octet
- Structure: S with 6 single bonds to F

20. N₂O₄ (Dinitrogen tetroxide)

- Each N: 5, each O: 6 → Total: 32 e⁻
- Two N atoms connected, each bonded to 2 O atoms
- Resonance structures with N-N single bond and double bonds to O
- Structure: O₂N-NO₂ with resonance

21. C₂O₄²⁻ (Oxalate ion)

- Each C: 4, each O: 6, +2 for negative charge → Total: 32 e⁻
- Two C atoms connected, each bonded to 2 O atoms
- Resonance structures with double bonds
- Structure: O=C-O⁻ ↔ O⁻-C=O with resonance

22. H₂PO₄⁻ (Dihydrogen phosphate)

- P: 5, each O: 6, each H: 1, +1 for negative charge → Total: 32 e⁻
- P in center, bonded to 4 O atoms (2 OH groups, 2 O atoms)
- P=O double bond likely
- Structure: P with 4 bonds to O, two O-H bonds

23. XeO₄ (Xenon tetroxide)

- Xe: 8, each O: 6 → Total: 32 e⁻
- Xe in center, bonded to 4 O atoms
- Likely Xe=O double bonds
- Structure: Xe with 4 double bonds to O

24. PO₃⁻ (Phosphite ion)

- P: 5, each O: 6, +1 for negative charge → Total: 24 e⁻
- P in center, bonded to 3 O atoms
- P-O single bonds, P has lone pair
- Structure: P with 3 single bonds to O, 1 lone pair

25. NO (Nitric oxide)

- N: 5, O: 6 → Total: 11 e⁻ (odd number)
- N-O double bond
- One unpaired electron on N or O
- Structure: N=O with one unpaired electron

Key Concepts:
- Formal charge = Valence e⁻ - (nonbonding e⁻ + ½ bonding e⁻)
- Resonance occurs when multiple valid Lewis structures exist
- Octet rule applies to most elements, but exceptions exist (expanded octets for period 3+ elements)
- Electronegativity determines where formal charges are placed

For molecular geometry, we'll need to apply VSEPR theory after completing these Lewis structures.