Worksheet 1 Molecular Geometry Polarity For each of the molecules ... - Free Printable

Let's go through each molecule or ion step by step, following the instructions:

1. Draw the Lewis dot structure (using the octet rule).
2. Predict molecular geometry using VSEPR theory.
3. Determine molecular polarity (based on symmetry and bond dipoles).
4. Identify intermolecular forces.

---

1. CH₃Br (Bromomethane)

- Lewis Structure:
- C is central atom bonded to 3 H atoms and 1 Br atom.
- C has 4 bonds (no lone pairs), Br has 3 lone pairs.
- All atoms satisfy octet rule (H: 2 electrons).
- Structure:
```
H
|
H--C--Br
|
H
```
With 3 lone pairs on Br.

- Molecular Geometry: Tetrahedral (4 bonding pairs, 0 lone pairs on C)

- Polarity: Polar – C–Br bond is polar; asymmetrical tetrahedral shape → net dipole.

- Intermolecular Forces: Dipole-dipole, London dispersion forces

---

2. N₃⁻ (Azide ion)

- Lewis Structure:
- Central N bonded to two terminal N atoms.
- Resonance structures: [N=N⁺=N⁻] ↔ [N⁻=N⁺=N]
- Total valence electrons: 5×3 + 1 = 16 e⁻
- Best resonance: linear with double bonds and formal charges.
- Structure:
```
[⁻N≡N⁺=N⁻] ↔ [N⁻=N⁺≡N]
```
But most stable: linear, with one triple bond and one single bond (with charge separation).

- Molecular Geometry: Linear (2 bonding domains, no lone pairs on central N)

- Polarity: Nonpolar – symmetric linear structure; dipole moments cancel.

- Intermolecular Forces: Ionic (as ion), but in solution: ion-dipole, or if neutralized: dipole-dipole & dispersion

> Note: N₃⁻ is an ion, so it doesn’t have typical IMF like molecules. In solid state: ionic lattice. In solution: ion-dipole.

---

3. H₂S (Hydrogen sulfide)

- Lewis Structure:
- S is central, bonded to two H atoms.
- S has 2 lone pairs.
- Total valence electrons: 6 + 2(1) = 8
- Structure:
```
H
|
S:
|
H
```
With two lone pairs on S.

- Molecular Geometry: Bent (or angular) – 2 bonding pairs, 2 lone pairs → electron geometry: tetrahedral

- Polarity: Polar – bent shape leads to net dipole moment.

- Intermolecular Forces: Dipole-dipole, London dispersion forces

---

4. CS₂ (Carbon disulfide)

- Lewis Structure:
- C is central, double bonded to two S atoms.
- C has 4 bonds (two double bonds), each S has 2 lone pairs.
- Structure: S=C=S

- Molecular Geometry: Linear – 2 bonding domains, no lone pairs on C

- Polarity: Nonpolar – symmetrical linear molecule; bond dipoles cancel.

- Intermolecular Forces: London dispersion forces only (nonpolar)

---

5. SO₂ (Sulfur dioxide)

- Lewis Structure:
- S is central, bonded to two O atoms.
- One double bond, one single bond with resonance.
- S has 1 lone pair.
- Total valence electrons: 6 + 2(6) = 18
- Resonance:
```
O=S-O ↔ O⁻-S⁺=O
```
With one lone pair on S.

- Molecular Geometry: Bent (trigonal planar electron geometry; 2 bonding pairs, 1 lone pair)

- Polarity: Polar – bent shape; bond dipoles don’t cancel.

- Intermolecular Forces: Dipole-dipole, London dispersion

---

6. CO₃²⁻ (Carbonate ion)

- Lewis Structure:
- C is central, bonded to three O atoms.
- Resonance: three equivalent structures with one double bond and two single bonds (with negative charges on O).
- Total valence electrons: 4 + 3(6) + 2 = 24
- Structure:
```
O⁻
|
O=C
\
O⁻
```
Resonance makes all bonds equal.

- Molecular Geometry: Trigonal planar (3 bonding domains, no lone pairs on C)

- Polarity: Nonpolar – symmetrical trigonal planar; dipoles cancel.

- Intermolecular Forces: Ion-dipole (if in solution), or ionic lattice (in solid); not a neutral molecule.

---

7. CCl₂O (Phosgene)

- Lewis Structure:
- C is central, double bonded to O and single bonded to two Cl atoms.
- C has 4 bonds (double bond counts as 2).
- Each Cl has 3 lone pairs, O has 2 lone pairs.
- Structure:
```
Cl
|
C=O
|
Cl
```

- Molecular Geometry: Trigonal planar (3 bonding domains, no lone pairs on C)

- Polarity: Polar – C=O bond is highly polar; Cl atoms are less electronegative than O, but asymmetry leads to net dipole.

- Intermolecular Forces: Dipole-dipole, London dispersion

---

8. OF₂ (Oxygen difluoride)

- Lewis Structure:
- O is central, bonded to two F atoms.
- O has 2 lone pairs.
- Total valence electrons: 6 + 2(7) = 20
- Structure:
```
F
|
O:
|
F
```
With two lone pairs on O.

- Molecular Geometry: Bent (tetrahedral electron geometry; 2 bonding pairs, 2 lone pairs)

- Polarity: Polar – bent shape; bond dipoles don’t cancel.

- Intermolecular Forces: Dipole-dipole, London dispersion

---

9. AsCl₃ (Arsenic trichloride)

- Lewis Structure:
- As is central, bonded to 3 Cl atoms.
- As has 1 lone pair.
- Total valence electrons: 5 + 3(7) = 26
- Structure:
```
Cl
|
As:
|
Cl
|
Cl
```
With one lone pair on As.

- Molecular Geometry: Trigonal pyramidal (tetrahedral electron geometry; 3 bonding pairs, 1 lone pair)

- Polarity: Polar – asymmetric shape; dipole moments don't cancel.

- Intermolecular Forces: Dipole-dipole, London dispersion

---

10. NH₄⁺ (Ammonium ion)

- Lewis Structure:
- N is central, bonded to 4 H atoms.
- N has positive charge; no lone pairs.
- Total valence electrons: 5 + 4(1) - 1 = 8
- Structure:
```
H
|
H--N⁺--H
|
H
```
Tetrahedral.

- Molecular Geometry: Tetrahedral (4 bonding pairs, 0 lone pairs)

- Polarity: Nonpolar – symmetrical; but ion, so not truly "nonpolar" in molecular sense. However, due to symmetry, dipole moment is zero.

- Intermolecular Forces: Ion-dipole (in solution), or ionic interactions (in solid); not a neutral molecule.

---

## ✔ Final Table:

| Molecule or ion | Electron Dot Structure | Molecular Geometry | Molecular Polarity | Intermolecular Forces |
|-----------------|------------------------|--------------------|---------------------|------------------------|
| CH₃Br | C bonded to 3H and Br; Br has 3 lone pairs | Tetrahedral | Polar | Dipole-dipole, London dispersion |
| N₃⁻ | Linear, resonance: [N≡N⁺–N⁻] ↔ [N⁻–N⁺≡N] | Linear | Nonpolar | Ion-dipole, ionic |
| H₂S | S bonded to 2H, 2 lone pairs on S | Bent | Polar | Dipole-dipole, London dispersion |
| CS₂ | S=C=S, each S has 2 lone pairs | Linear | Nonpolar | London dispersion |
| SO₂ | Bent with resonance; S has 1 lone pair | Bent | Polar | Dipole-dipole, London dispersion |
| CO₃²⁻ | Trigonal planar with resonance | Trigonal planar | Nonpolar | Ion-dipole, ionic |
| CCl₂O | C bonded to O (double) and two Cl | Trigonal planar | Polar | Dipole-dipole, London dispersion |
| OF₂ | O bonded to 2F, 2 lone pairs on O | Bent | Polar | Dipole-dipole, London dispersion |
| AsCl₃ | As bonded to 3Cl, 1 lone pair on As | Trigonal pyramidal | Polar | Dipole-dipole, London dispersion |
| NH₄⁺ | N bonded to 4H, no lone pairs, +1 charge | Tetrahedral | Nonpolar (symmetric) | Ion-dipole, ionic |

---

🔍 Summary Notes:

- Nonpolar molecules: Symmetrical shape and/or identical substituents cancel dipoles.
- Polar molecules: Asymmetrical shape or unequal electronegativity causes net dipole.
- Ions (like N₃⁻, CO₃²⁻, NH₄⁺): Exhibit ionic behavior or ion-dipole forces rather than typical IMFs.
- Intermolecular forces depend on polarity and molecular type:
- London dispersion: All molecules (especially nonpolar)
- Dipole-dipole: Polar molecules
- Ion-dipole / ionic: Ions in solution or solids

Let me know if you'd like hand-drawn Lewis structures or diagrams!