1️⃣ Quick Review: Molecules and Compounds

✅ Molecules

A molecule is made of two or more atoms bonded together.
Examples: O₂, H₂, H₂O

✅ Compounds (Heteronuclear Molecules)

A compound is a molecule made of different kinds of atoms.
Examples: H₂O, CO₂, NH₃, NaCl

📌 Why this matters:
Polarity only makes sense when atoms are bonded — especially in covalent molecules.

2️⃣ Review: Covalent vs Ionic Bonding

✅ Covalent Bonding

• atoms share electrons

• usually between nonmetals

• sharing may be equal or unequal
  
  

✅ Ionic Bonding

• electrons are transferred

• usually metal + nonmetal

• creates charged ions
  
  

📌 Polarity is mainly studied in covalent molecules because sharing can be uneven.

3️⃣ Bond Polarity and Electronegativity

✅ Electronegativity (EN)

Electronegativity is an atom’s ability to pull shared electrons toward itself.

• higher EN → stronger pull

• lower EN → weaker pull
  
  

✅ Bond Polarity

A covalent bond is:

• Nonpolar covalent if electrons are shared equally

• Polar covalent if electrons are shared unequally
  
  

📌 Electronegativity Difference (ΔEN):

• 0.0–0.4 → Nonpolar covalent

• 0.5–1.7 → Polar covalent

• 1.8+ → Ionic
  
  

Partial Charges

In polar bonds:

• atom with higher EN becomes δ⁻

• atom with lower EN becomes δ⁺
  
  

Example: H–Cl
Cl pulls harder → Cl is δ⁻, H is δ⁺

4️⃣ Lewis Structures

✅ What is a Lewis Structure?

A Lewis structure shows:

• valence electrons

• bonds (shared pairs)

• lone pairs
  
  

It helps you identify:
✅ bond polarity
✅ molecular shape
✅ molecular polarity

🪄 Steps to Draw Lewis Structures

1. Count total valence electrons.

2. Choose the central atom (least EN, never H).

3. Connect atoms with single bonds.

4. Complete octets of outside atoms.

5. Place remaining electrons on the central atom.

6. If needed, form double/triple bonds.

✏️ Lewis Structure Examples (Core Models)

H₂O (Water)

Total valence e⁻ = 8
O central, 2 bonds, 2 lone pairs.

CO₂ (Carbon Dioxide)

Total valence e⁻ = 16
C central, 2 double bonds.

NH₃ (Ammonia)

Total valence e⁻ = 8
N central, 3 bonds, 1 lone pair.

CH₄ (Methane)

Total valence e⁻ = 8
C central, 4 bonds, no lone pairs.

5️⃣ Molecular Geometry (VSEPR – Basic Intro)

✅ Why shape matters

Even if bonds are polar, the whole molecule may be polar or nonpolar depending on shape.

✅ VSEPR idea

Electron pairs repel → spread out → shape forms.

Basic shapes you need:

6️⃣ Determining Molecular Polarity

✅ Step 1: Check bond polarity

Are the bonds polar? (ΔEN check)

✅ Step 2: Check shape + symmetry

A. Symmetrical molecules → NONPOLAR

Polar bonds cancel out.

Net dipole moment = 0

Example: CO₂, CH₄

B. Asymmetrical molecules → POLAR

Polar bonds don’t cancel.

Net dipole moment ≠ 0

Example: H₂O, NH₃