3.3 Polar and Nonpolar Covalent Bond

This topic explains the difference between polar and nonpolar covalent bonds, how they arise from electronegativity differences, and how molecular shape determines whether a molecule is polar or nonpolar.

Definition

Polar Covalent Bond: A covalent bond in which the shared electrons are unequally distributed because of a difference in electronegativity between the bonded atoms.

Nonpolar Covalent Bond: A covalent bond in which the shared electrons are equally distributed because the bonded atoms have the same or very similar electronegativity.

Main Idea

In a polar covalent bond, the more electronegative atom attracts the shared electrons toward itself, acquiring a partial negative charge (δ−), while the less electronegative atom acquires a partial positive charge (δ+). Such molecules are called polar molecules and have a dipole moment.

In a nonpolar covalent bond, the electrons are shared equally, and there is no separation of charge.

Core Concept

The polarity of a bond depends on the electronegativity difference between the bonded atoms. However, the overall polarity of a molecule depends not only on bond polarity but also on the shape of the molecule. Symmetrical molecules can have polar bonds but still be nonpolar overall if the bond dipoles cancel out.

Detailed Explanation

Polar Covalent Bond

• In a polar covalent bond, due to electronegativity difference, one atom (which is more electronegative) will attract the shared electron to itself and bear a partial negative charge (δ−).
• The other atom bears a partial positive charge (δ+).
• Such molecules are called polar molecules.
• Polar molecules have dipoles — electric charges of equal magnitude and opposite sign.

Nonpolar Covalent Bond

• In a nonpolar covalent bond, the electrons are shared equally.
• There is no separation of charge.
• Nonpolar molecules have no overall charge and no dipole moment.

Symmetrical Polyatomic Molecules

• Symmetrical polyatomic molecules are nonpolar molecules.
• They have shapes such as:
• Linear
• Trigonal Planar
• Tetrahedral
• In these molecules, the dipole of each bond exerts equal and opposite effects, and hence the dipoles cancel each other out.
• Nonpolar molecules have no overall charge.

Example: Carbon Dioxide (CO₂)

• Due to the electronegativity difference between carbon and oxygen, the C=O bond in CO₂ is polar.
• However, CO₂ is nonpolar because the two bond dipoles are equal and opposite (the molecule is linear), so they cancel each other out.
• The overall dipole moment of CO₂ is zero.

Summary Table

Concept Assessment Exercise 3.1

Are the following molecules polar or non-polar? In each case, give a reason for your answer.

(Electronegativity values: F = 4.0, Cl = 3.0, Br = 2.8, S = 2.5, C = 2.5, H = 2.1)

• Both atoms are Cl (EN = 3.0).
• Difference = 0.
• Electrons are shared equally.
• Nonpolar.

• H (EN = 2.1), F (EN = 4.0).
• Difference = 1.9.
• F is more electronegative → δ−; H → δ+.
• Polar.

• S (EN = 2.5), Cl (EN = 3.0).
• Difference = 0.5.
• S is less electronegative → δ+; Cl → δ−.
• The molecule is V-shaped, so the bond dipoles do not cancel.
• Polar.

• C (EN = 2.5), H (EN = 2.1), Cl (EN = 3.0).
• C−H difference = 0.4; C−Cl difference = 0.5.
• The C−Cl bond is more polar than the C−H bonds.
• The molecule is tetrahedral, and the dipoles do not cancel.
• Polar.

• C (EN = 2.5), Br (EN = 2.8).
• Difference = 0.3 (small, but polar bonds).
• The molecule is tetrahedral and symmetrical.
• The bond dipoles cancel each other out.
• Nonpolar.

Intermediate Level

Polar Covalent Bond

• Unequal sharing of electrons.
• Due to electronegativity difference.
• Creates partial charges: δ+ and δ−.
• Example: HCl, HF, H₂O.

Nonpolar Covalent Bond

• Equal sharing of electrons.
• No electronegativity difference (or very small).
• No partial charges.
• Example: Cl₂, H₂, O₂.

Molecular Polarity

• Depends on both bond polarity and molecular shape.
• Symmetrical molecules → dipoles cancel → nonpolar.
• Asymmetrical molecules → dipoles do not cancel → polar.

Examples

Important Points

• 1. Polar covalent bond: Unequal sharing of electrons; δ+ and δ−.
• 2. Nonpolar covalent bond: Equal sharing of electrons; no partial charges.
• 3. Molecular polarity: Depends on bond polarity and molecular shape.
• 4. Symmetrical molecules: Dipoles cancel → nonpolar.
• 5. Asymmetrical molecules: Dipoles do not cancel → polar.
• 6. Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar), CHCl₃ (polar).

Remember

• - Polar covalent: Unequal sharing; δ+ and δ−
• - Nonpolar covalent: Equal sharing; no partial charges
• - Symmetrical: Dipoles cancel → nonpolar
• - Asymmetrical: Dipoles do not cancel → polar
• - Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar)

Common Mistakes

• 1. Thinking all polar bonds make a polar molecule: Symmetrical molecules can have zero dipole moment (e.g., CO₂, CCl₄).
• 2. Forgetting the role of molecular shape: Shape determines whether dipoles cancel.
• 3. Confusing bond polarity with molecular polarity: Bond polarity is about individual bonds; molecular polarity is about the whole molecule.
• 4. Ignoring small electronegativity differences: Even small differences (0.4–1.8) make a bond polar.

Quick Revision

• Polar covalent bond: Unequal sharing of electrons; δ+ and δ−.
• Nonpolar covalent bond: Equal sharing of electrons; no partial charges.
• Molecular polarity depends on bond polarity and molecular shape.
• Symmetrical molecules: Dipoles cancel → nonpolar.
• Asymmetrical molecules: Dipoles do not cancel → polar.
• Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar), CHCl₃ (polar).

60-Second Revision

• Polar covalent: Unequal sharing; δ+ and δ−
• Nonpolar covalent: Equal sharing; no partial charges
• Symmetrical: Dipoles cancel → nonpolar
• Asymmetrical: Dipoles do not cancel → polar
• Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar)

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