Worksheet for practicing Lewis structures of various chemical compounds with instructions to include resonance structures and determine preferred structures using formal charges.

The task involves drawing Lewis structures for a variety of molecules and ions, including resonance structures where applicable. Additionally, formal charges should be used to determine the most preferred Lewis structure when multiple options exist. Below, I will outline the general steps for solving these problems and provide examples for a few molecules/ions.

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General Steps for Drawing Lewis Structures

1. Calculate the Total Number of Valence Electrons:
- Sum the valence electrons from all atoms.
- Add extra electrons for negative charges or subtract electrons for positive charges.

2. Determine the Central Atom:
- The central atom is typically the least electronegative atom (except for hydrogen, which is never central).
- For polyatomic ions, the central atom is often the one with the highest bonding capacity.

3. Draw a Skeletal Structure:
- Connect atoms with single bonds to form a skeleton.
- Ensure that each atom has an octet (or duet for hydrogen).

4. Place Remaining Electrons as Lone Pairs:
- Distribute remaining electrons as lone pairs on terminal atoms first, then on the central atom.

5. Check for Octets/Duets:
- Ensure all atoms have complete valence shells (octets or duets).

6. Form Double or Triple Bonds if Necessary:
- If the central atom does not have an octet, use lone pairs to form double or triple bonds.

7. Identify Resonance Structures:
- Draw alternative Lewis structures by moving electrons (but not atoms) to satisfy the octet rule.
- Use formal charges to determine the most stable structure.

8. Calculate Formal Charges:
- Formal charge = [Valence electrons] - [Non-bonding electrons] - 0.5 × [Bonding electrons].
- The most stable structure has the smallest formal charges and charges closest to zero.

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Examples of Lewis Structures

#### Example 1: CH₄ (Methane)
1. Total Valence Electrons:
- Carbon (C): 4
- Hydrogen (H): 1 × 4 = 4
- Total: 4 + 4 = 8

2. Skeletal Structure:
- C is the central atom, bonded to 4 H atoms.

3. Lewis Structure:
```
H
|
H—C—H
|
H
```

4. Resonance Structures:
- No resonance structures are possible.

5. Formal Charges:
- C: 4 - 0 - 8/2 = 0
- H: 1 - 0 - 2/2 = 0
- All formal charges are zero, so this is the correct structure.

---

#### Example 2: BF₃ (Boron Trifluoride)
1. Total Valence Electrons:
- Boron (B): 3
- Fluorine (F): 7 × 3 = 21
- Total: 3 + 21 = 24

2. Skeletal Structure:
- B is the central atom, bonded to 3 F atoms.

3. Lewis Structure:
```
F
|
F—B—F
```

4. Resonance Structures:
- No resonance structures are possible.

5. Formal Charges:
- B: 3 - 0 - 6/2 = 0
- F: 7 - 6 - 2/2 = 0
- All formal charges are zero, so this is the correct structure.

---

#### Example 3: N₃⁻ (Azide Ion)
1. Total Valence Electrons:
- Nitrogen (N): 5 × 3 = 15
- Charge: +1 electron
- Total: 15 + 1 = 16

2. Skeletal Structure:
- Linear arrangement: N—N—N

3. Lewis Structure:
```
:N≡N⁺−N:
```

4. Resonance Structures:
- There are three resonance structures:
```
Structure 1: :N≡N⁺−N:
Structure 2: :N−N=N⁺:
Structure 3: :N=N⁺−N:
```

5. Formal Charges:
- In all structures, the central N atom has a +1 charge, and the terminal N atoms have a −1 charge. These are the most stable resonance structures.

---

#### Example 4: H₂O (Water)
1. Total Valence Electrons:
- Oxygen (O): 6
- Hydrogen (H): 1 × 2 = 2
- Total: 6 + 2 = 8

2. Skeletal Structure:
- O is the central atom, bonded to 2 H atoms.

3. Lewis Structure:
```
H
|
H—O:
```

4. Resonance Structures:
- No resonance structures are possible.

5. Formal Charges:
- O: 6 - 4 - 4/2 = 0
- H: 1 - 0 - 2/2 = 0
- All formal charges are zero, so this is the correct structure.

---

#### Example 5: NO₂⁻ (Nitrite Ion)
1. Total Valence Electrons:
- Nitrogen (N): 5
- Oxygen (O): 6 × 2 = 12
- Charge: +1 electron
- Total: 5 + 12 + 1 = 18

2. Skeletal Structure:
- N is the central atom, bonded to 2 O atoms.

3. Lewis Structure:
```
:O=N−O:
```

4. Resonance Structures:
- There are two resonance structures:
```
Structure 1: :O=N−O:
Structure 2: :O−N=O:
```

5. Formal Charges:
- In both structures, the central N atom has a +1 charge in one structure and a −1 charge in the other. The structure with the more electronegative atom (O) having the negative charge is preferred.

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Summary

- For each molecule/ion, follow the steps above to draw the Lewis structure.
- Identify resonance structures when possible and use formal charges to determine the most stable structure.
- Practice these steps for all the given molecules/ions in the list.

If you need detailed solutions for specific molecules, feel free to ask!

Final Answer:
\boxed{\text{Follow the steps outlined above to solve for each molecule/ion.}}