SOLUTION: CHEM 10 SMC Lewis Structures and Molecular Shapes ... - Free Printable

Problem Analysis and Solution Explanation

The task involves analyzing the Lewis structures, molecular shapes, bond angles, and polarity of several molecules. Below is a detailed explanation for each molecule:

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#### 1. H₂S (Hydrogen Sulfide)

- Total # of Valence Electrons:
- Hydrogen (H): 1 valence electron × 2 = 2
- Sulfur (S): 6 valence electrons
- Total: \( 2 + 6 = 8 \) valence electrons

- Lewis Structure:
- Sulfur (S) is the central atom.
- Each hydrogen (H) forms a single bond with sulfur.
- Sulfur has 2 lone pairs of electrons to complete its octet.
- Structure:
```
H—S—H
:
```

- 3-D Model Sketch:
- The molecule is bent due to the presence of two lone pairs on sulfur.

- Bond Angles:
- The ideal bond angle for a tetrahedral arrangement is \( 109.5^\circ \), but the lone pairs cause slight compression, so the actual bond angle is slightly less than \( 109.5^\circ \).

- Molecular Shape:
- Bent (V-shaped)

- Polar Bonds:
- Yes, because the S-H bonds are polar (different electronegativities).

- Molecular Polarity:
- Polar, due to the bent shape and the presence of polar bonds, which results in a net dipole moment.

---

#### 2. COCl₂ (Carbon Dioxide Chloride or Phosgene)

- Total # of Valence Electrons:
- Carbon (C): 4 valence electrons
- Oxygen (O): 6 valence electrons
- Chlorine (Cl): 7 valence electrons × 2 = 14
- Total: \( 4 + 6 + 14 = 24 \) valence electrons

- Lewis Structure:
- Carbon (C) is the central atom.
- Carbon forms double bonds with oxygen and single bonds with each chlorine atom.
- Oxygen has 2 lone pairs, and each chlorine has 3 lone pairs.
- Structure:
```
:O::C::Cl:
|
Cl:
```

- 3-D Model Sketch:
- The molecule is trigonal planar around the central carbon atom.

- Bond Angles:
- Ideal bond angle for trigonal planar geometry: \( 120^\circ \).

- Molecular Shape:
- Trigonal planar

- Polar Bonds:
- Yes, because the C-O and C-Cl bonds are polar (different electronegativities).

- Molecular Polarity:
- Polar, due to the presence of polar bonds and the asymmetry of the molecule, resulting in a net dipole moment.

---

#### 3. SiI₄ (Silicon Tetraiodide)

- Total # of Valence Electrons:
- Silicon (Si): 4 valence electrons
- Iodine (I): 7 valence electrons × 4 = 28
- Total: \( 4 + 28 = 32 \) valence electrons

- Lewis Structure:
- Silicon (Si) is the central atom.
- Silicon forms single bonds with each iodine atom.
- Each iodine has 3 lone pairs to complete its octet.
- Structure:
```
:I—I—Si—I—I:
```

- 3-D Model Sketch:
- The molecule is tetrahedral around the central silicon atom.

- Bond Angles:
- Ideal bond angle for tetrahedral geometry: \( 109.5^\circ \).

- Molecular Shape:
- Tetrahedral

- Polar Bonds:
- Yes, because the Si-I bonds are polar (different electronegativities).

- Molecular Polarity:
- Non-polar, because the tetrahedral symmetry cancels out the individual bond dipoles, resulting in no net dipole moment.

---

#### 4. HCN (Hydrogen Cyanide)

- Total # of Valence Electrons:
- Hydrogen (H): 1 valence electron
- Carbon (C): 4 valence electrons
- Nitrogen (N): 5 valence electrons
- Total: \( 1 + 4 + 5 = 10 \) valence electrons

- Lewis Structure:
- Carbon (C) is the central atom.
- Carbon forms a triple bond with nitrogen and a single bond with hydrogen.
- Nitrogen has 1 lone pair.
- Structure:
```
H—C≡N:
```

- 3-D Model Sketch:
- The molecule is linear.

- Bond Angles:
- Ideal bond angle for a linear molecule: \( 180^\circ \).

- Molecular Shape:
- Linear

- Polar Bonds:
- Yes, because the C-N and C-H bonds are polar (different electronegativities).

- Molecular Polarity:
- Polar, due to the linear shape and the presence of polar bonds, which result in a net dipole moment.

---

Final Answers

1. H₂S:
- Molecular Shape: Bent
- Bond Angles: \( 109.5^\circ \)
- Molecular Polarity: Polar

2. COCl₂:
- Molecular Shape: Trigonal Planar
- Bond Angles: \( 120^\circ \)
- Molecular Polarity: Polar

3. SiI₄:
- Molecular Shape: Tetrahedral
- Bond Angles: \( 109.5^\circ \)
- Molecular Polarity: Non-Polar

4. HCN:
- Molecular Shape: Linear
- Bond Angles: \( 180^\circ \)
- Molecular Polarity: Polar

Boxed Final Answer:

\[
\boxed{
\begin{array}{l}
\text{1. H₂S: Bent, } 109.5^\circ, \text{ Polar} \\
\text{2. COCl₂: Trigonal Planar, } 120^\circ, \text{ Polar} \\
\text{3. SiI₄: Tetrahedral, } 109.5^\circ, \text{ Non-Polar} \\
\text{4. HCN: Linear, } 180^\circ, \text{ Polar}
\end{array}
}
\]