Lewis Structure Worksheet | Lecture notes Geometry | Docsity - Free Printable

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 is a detailed explanation of how to approach this problem.

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Steps to Solve the Problem

1. Determine the Total Number of Valence Electrons:
- For neutral molecules: Sum the valence electrons of all atoms.
- For ions: Add or subtract the charge (e.g., add 1 electron for a negative ion, subtract 1 electron for a positive ion).

2. Choose 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 that can expand its octet (e.g., sulfur, phosphorus, chlorine).

3. Draw a Skeletal Structure:
- Arrange the atoms around the central atom.
- Connect them with single bonds initially.

4. Distribute Remaining Electrons:
- Place lone pairs on terminal atoms first to satisfy the octet rule.
- If the central atom does not have an octet, use multiple bonds (double or triple) to complete its octet.

5. Check for Resonance Structures:
- Resonance occurs when there are multiple valid Lewis structures due to the movement of electrons.
- Draw all possible resonance structures and use formal charges to determine the most stable one.

6. Calculate Formal Charges:
- Formal charge = [Valence electrons] - [Non-bonding electrons] - [Bonding electrons / 2].
- The most stable structure has the smallest formal charges, ideally zero or close to zero.

7. Verify the Octet Rule:
- Ensure that all atoms (except hydrogen) have 8 electrons in their valence shell (octet rule). Exceptions include:
- Hydrogen: 2 electrons.
- Elements in Period 3 or higher: Can exceed the octet rule (e.g., sulfur, phosphorus).

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Example Solutions

#### 1. CH₄ (Methane)
- Total valence electrons: \(4 + 4 \times 1 = 8\)
- Central atom: Carbon (C)
- Skeletal structure: C-H bonds
- Final structure:
```
H
|
H-C-H
|
H
```
- Resonance: None.
- Formal charges: All atoms have a formal charge of 0.

#### 2. BF₃ (Boron Trifluoride)
- Total valence electrons: \(3 + 3 \times 7 = 24\)
- Central atom: Boron (B)
- Skeletal structure: B-F bonds
- Final structure:
```
F
|
F-B-F
```
- Resonance: None.
- Formal charges: All atoms have a formal charge of 0.

#### 3. HCN (Hydrogen Cyanide)
- Total valence electrons: \(1 + 4 + 5 = 10\)
- Central atom: Carbon (C)
- Skeletal structure: H-C-N
- Final structure:
```
:N≡C:H
```
- Resonance: None.
- Formal charges: All atoms have a formal charge of 0.

#### 4. NH₃ (Ammonia)
- Total valence electrons: \(5 + 3 \times 1 = 8\)
- Central atom: Nitrogen (N)
- Skeletal structure: N-H bonds
- Final structure:
```
H
|
H-N-H
|
H
```
- Resonance: None.
- Formal charges: All atoms have a formal charge of 0.

#### 5. H₂O (Water)
- Total valence electrons: \(6 + 2 \times 1 = 8\)
- Central atom: Oxygen (O)
- Skeletal structure: O-H bonds
- Final structure:
```
H
|
H-O
|
H
```
- Resonance: None.
- Formal charges: All atoms have a formal charge of 0.

#### 6. ClO₃⁻ (Chlorate Ion)
- Total valence electrons: \(7 + 3 \times 6 + 1 = 26\)
- Central atom: Chlorine (Cl)
- Skeletal structure: Cl-O bonds
- Final structure:
```
O
|
O-Cl-O
|
O
```
- Resonance: Yes. There are three equivalent resonance structures.
- Formal charges: In the most stable structure, Cl has a formal charge of +1, and one oxygen has a formal charge of -1.

#### 7. H₃O⁺ (Hydronium Ion)
- Total valence electrons: \(6 + 3 \times 1 - 1 = 8\)
- Central atom: Oxygen (O)
- Skeletal structure: O-H bonds
- Final structure:
```
H
|
H-O-H
|
H
```
- Resonance: Yes. There are three equivalent resonance structures.
- Formal charges: In the most stable structure, O has a formal charge of +1, and one hydrogen has a formal charge of -1.

#### 8. NH₂⁻ (Amide Ion)
- Total valence electrons: \(5 + 2 \times 1 + 1 = 8\)
- Central atom: Nitrogen (N)
- Skeletal structure: N-H bonds
- Final structure:
```
H
|
H-N:
|
H
```
- Resonance: Yes. There are two equivalent resonance structures.
- Formal charges: In the most stable structure, N has a formal charge of -1, and one hydrogen has a formal charge of +1.

#### 9. NH₄⁺ (Ammonium Ion)
- Total valence electrons: \(5 + 4 \times 1 - 1 = 8\)
- Central atom: Nitrogen (N)
- Skeletal structure: N-H bonds
- Final structure:
```
H
|
H-N-H
|
H
```
- Resonance: No.
- Formal charges: All atoms have a formal charge of 0.

#### 10. C₂H₄ (Ethene)
- Total valence electrons: \(2 \times 4 + 4 \times 1 = 12\)
- Central atoms: Two carbon atoms (C=C double bond)
- Skeletal structure: C=C with H atoms
- Final structure:
```
H
|
H-C=C-H
|
H
```
- Resonance: No.
- Formal charges: All atoms have a formal charge of 0.

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General Approach for All Molecules/Ions

- Follow the steps above for each molecule/ion.
- Pay special attention to:
- Ions: Adjust the total number of valence electrons based on the charge.
- Resonance: Draw all possible resonance structures and use formal charges to determine the most stable one.
- Expanded octets: Elements like sulfur, phosphorus, and chlorine can exceed the octet rule.

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Final Answer

The solutions involve drawing Lewis structures for each molecule/ion, considering resonance and formal charges as necessary. The final answer is:

\boxed{\text{See detailed explanations and structures above.}}