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, 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 that can accommodate more than an octet.

3. Draw a Skeletal Structure:
- Connect atoms with single bonds to form a skeleton.
- Ensure each atom has at least one bond unless it is a noble gas.

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 Formal Charges:
- Calculate formal charges for each atom:
\[
\text{Formal Charge} = \text{Valence Electrons} - \text{Non-bonding Electrons} - \frac{\text{Bonding Electrons}}{2}
\]
- The best Lewis structure minimizes formal charges and keeps them as close to zero as possible.

6. Identify Resonance Structures:
- Draw alternative Lewis structures if electrons can be rearranged while maintaining the same connectivity.
- Use formal charges to determine the most stable resonance structure.

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Examples of Lewis Structures

#### Example 1: CH₄ (Methane)
1. Total Valence Electrons:
\( C \): 4 valence electrons
\( H \): 1 valence electron × 4 = 4 valence electrons
Total = 4 + 4 = 8

2. Skeletal Structure:
Carbon is the central atom, bonded to four hydrogens.
\[
\ce{H-C-H}
\]
\[
\ce{ | | }
\]
\[
\ce{H H}
\]

3. Lone Pairs:
All electrons are used in bonding, so no lone pairs.

4. Formal Charges:
- Carbon: \( 4 - 0 - \frac{8}{2} = 0 \)
- Hydrogen: \( 1 - 0 - \frac{2}{2} = 0 \)

All formal charges are zero, so this is the correct structure.

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

2. Skeletal Structure:
Boron is the central atom, bonded to three fluorines.
\[
\ce{F-B-F}
\]
\[
\ce{ | | }
\]
\[
\ce{F}
\]

3. Lone Pairs:
Each fluorine has three lone pairs. Boron has no lone pairs.

4. Formal Charges:
- Boron: \( 3 - 0 - \frac{6}{2} = 0 \)
- Fluorine: \( 7 - 6 - \frac{2}{2} = 0 \)

All formal charges are zero, so this is the correct structure.

#### Example 3: NO₂⁻ (Nitrite Ion)
1. Total Valence Electrons:
\( N \): 5 valence electrons
\( O \): 6 valence electrons × 2 = 12 valence electrons
Extra electron for the negative charge = 1
Total = 5 + 12 + 1 = 18

2. Skeletal Structure:
Nitrogen is the central atom, bonded to two oxygens.
\[
\ce{O-N-O}
\]

3. Lone Pairs and Resonance:
- Each oxygen needs 6 electrons (including the bond), so place lone pairs accordingly.
- There are two resonance structures:
\[
\ce{O=N-O^-} \quad \text{and} \quad \ce{O^-=N-O}
\]

4. Formal Charges:
- For \(\ce{O=N-O^-}\):
- Nitrogen: \( 5 - 0 - \frac{4}{2} = +1 \)
- Oxygen (double bond): \( 6 - 4 - \frac{4}{2} = 0 \)
- Oxygen (single bond): \( 6 - 6 - \frac{2}{2} = -1 \)
- For \(\ce{O^-=N-O}\):
- Nitrogen: \( 5 - 0 - \frac{4}{2} = +1 \)
- Oxygen (double bond): \( 6 - 4 - \frac{4}{2} = 0 \)
- Oxygen (single bond): \( 6 - 6 - \frac{2}{2} = -1 \)

Both resonance structures have the same formal charges, so they are equally preferred.

#### Example 4: NH₄⁺ (Ammonium Ion)
1. Total Valence Electrons:
\( N \): 5 valence electrons
\( H \): 1 valence electron × 4 = 4 valence electrons
Subtract 1 electron for the positive charge = -1
Total = 5 + 4 - 1 = 8

2. Skeletal Structure:
Nitrogen is the central atom, bonded to four hydrogens.
\[
\ce{H-N-H}
\]
\[
\ce{ | | }
\]
\[
\ce{H H}
\]

3. Lone Pairs:
All electrons are used in bonding, so no lone pairs.

4. Formal Charges:
- Nitrogen: \( 5 - 0 - \frac{8}{2} = +1 \)
- Hydrogen: \( 1 - 0 - \frac{2}{2} = 0 \)

The formal charge on nitrogen matches the charge of the ion, so this is the correct structure.

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Summary

- For each molecule/ion, follow the steps above to draw the Lewis structure.
- Pay attention to formal charges to determine the most preferred structure.
- Identify resonance structures when applicable and use formal charges to evaluate their stability.

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

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