Worksheet 1: Molecular Geometry & Polarity - Students analyze molecules and ions by drawing Lewis structures, predicting molecular geometry, determining polarity, and identifying intermolecular forces.

To solve the problem, we need to analyze each molecule or ion step by step using the following guidelines:

1. Draw the Lewis Dot Structure: Ensure that all atoms follow the Octet Rule.
2. Predict Molecular Geometry: Use VSEPR (Valence Shell Electron Pair Repulsion) theory to determine the geometry.
3. Determine Molecular Polarity: Assess whether the molecule is polar or nonpolar based on its geometry and bond dipoles.
4. Identify Intermolecular Forces: Determine the types of intermolecular forces present.

Let's go through each molecule or ion:

---

1. CH₃Br

- Lewis Dot Structure: Carbon (C) is the central atom with single bonds to three hydrogen (H) atoms and one bromine (Br) atom. All atoms satisfy the Octet Rule.
```
H
|
H-C-Br
|
H
```
- Molecular Geometry: Tetrahedral (all four groups around C are bonding pairs).
- Molecular Polarity: Polar due to the difference in electronegativity between C and Br, resulting in a net dipole moment.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

2. Nl₃

- Lewis Dot Structure: Nitrogen (N) is the central atom with single bonds to three iodine (I) atoms. N has one lone pair.
```
I
|
I-N-I
|
I
```
- Molecular Geometry: Trigonal pyramidal (three bonding pairs and one lone pair around N).
- Molecular Polarity: Polar due to the lone pair and the asymmetry of the molecule.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

3. H₂S

- Lewis Dot Structure: Sulfur (S) is the central atom with single bonds to two hydrogen (H) atoms. S has two lone pairs.
```
H
|
H-S
|
H
```
- Molecular Geometry: Bent (two bonding pairs and two lone pairs around S).
- Molecular Polarity: Polar due to the lone pairs and the asymmetry of the molecule.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

4. CS₂

- Lewis Dot Structure: Carbon (C) is the central atom with double bonds to two sulfur (S) atoms. No lone pairs on C.
```
S
||
S=C=S
||
S
```
- Molecular Geometry: Linear (two double bonds around C).
- Molecular Polarity: Nonpolar because the molecule is symmetrical and the bond dipoles cancel out.
- Intermolecular Forces: London dispersion forces.

---

5. SO₂

- Lewis Dot Structure: Sulfur (S) is the central atom with double bonds to two oxygen (O) atoms and one lone pair.
```
O
||
O-S
|
:
```
- Molecular Geometry: Bent (two bonding pairs and one lone pair around S).
- Molecular Polarity: Polar due to the lone pair and the asymmetry of the molecule.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

6. CO₃²⁻

- Lewis Dot Structure: Carbon (C) is the central atom with double bonds to two oxygen (O) atoms and a single bond to one oxygen (O) atom. The negative charge is distributed among the three oxygens.
```
O
||
O-C-O
|
:
```
- Molecular Geometry: Trigonal planar (three bonding pairs around C).
- Molecular Polarity: Nonpolar because the molecule is symmetrical and the bond dipoles cancel out.
- Intermolecular Forces: Ion-dipole interactions (due to the negative charge).

---

7. CCl₂O

- Lewis Dot Structure: Carbon (C) is the central atom with single bonds to two chlorine (Cl) atoms and one oxygen (O) atom, and a double bond to another oxygen (O) atom.
```
Cl
|
Cl-C-O
|
O
```
- Molecular Geometry: Tetrahedral (four groups around C: two single bonds, one double bond, and one lone pair on O).
- Molecular Polarity: Polar due to the difference in electronegativity between C and Cl/O.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

8. OF₂

- Lewis Dot Structure: Oxygen (O) is the central atom with single bonds to two fluorine (F) atoms and two lone pairs.
```
F
|
F-O
|
:
```
- Molecular Geometry: Bent (two bonding pairs and two lone pairs around O).
- Molecular Polarity: Polar due to the lone pairs and the asymmetry of the molecule.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

9. AsCl₃

- Lewis Dot Structure: Arsenic (As) is the central atom with single bonds to three chlorine (Cl) atoms and one lone pair.
```
Cl
|
Cl-As-Cl
|
:
```
- Molecular Geometry: Trigonal pyramidal (three bonding pairs and one lone pair around As).
- Molecular Polarity: Polar due to the lone pair and the asymmetry of the molecule.
- Intermolecular Forces: Dipole-dipole interactions and London dispersion forces.

---

10. NH₄⁺

- Lewis Dot Structure: Nitrogen (N) is the central atom with single bonds to four hydrogen (H) atoms. The positive charge is due to the lack of a lone pair.
```
H
|
H-N-H
|
H
```
- Molecular Geometry: Tetrahedral (four bonding pairs around N).
- Molecular Polarity: Nonpolar because the molecule is symmetrical.
- Intermolecular Forces: Ion-dipole interactions (due to the positive charge).

---

Final Answer:

\[
\boxed{
\begin{array}{|c|c|c|c|c|}
\hline
\text{Molecule or ion} & \text{Lewis Dot Structure} & \text{Molecular Geometry} & \text{Molecular Polarity} & \text{Intermolecular Forces} \\
\hline
\text{CH}_3\text{Br} & \text{Tetrahedral} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{Nl}_3 & \text{Trigonal pyramidal} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{H}_2\text{S} & \text{Bent} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{CS}_2 & \text{Linear} & \text{Nonpolar} & \text{London} \\
\hline
\text{SO}_2 & \text{Bent} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{CO}_3^{2-} & \text{Trigonal planar} & \text{Nonpolar} & \text{Ion-dipole} \\
\hline
\text{CCl}_2\text{O} & \text{Tetrahedral} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{OF}_2 & \text{Bent} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{AsCl}_3 & \text{Trigonal pyramidal} & \text{Polar} & \text{Dipole-dipole, London} \\
\hline
\text{NH}_4^+ & \text{Tetrahedral} & \text{Nonpolar} & \text{Ion-dipole} \\
\hline
\end{array}
}
\]