3.2 Dipole Moment

This topic explains dipole moment — what it is, how it is calculated, and how it determines the polarity of a molecule.

Definition

Dipole moment is the product of the magnitude of positive or negative charges and the distance between them.

μ = q × d

Units

• SI unit: Debye (D)
• 1 Debye = 3.335 × 10^-30 C·m

Electronegativity Difference and Dipole Moment

• Greater electronegativity difference → more polar bond → greater dipole moment.
• Smaller electronegativity difference → less polar bond → smaller dipole moment.

Bond Polarity vs. Molecular Polarity

• HCl: Polar bond → polar molecule
• CO₂: Polar bonds → nonpolar molecule (dipoles cancel)

Representation

• Arrow from δ+ to δ−
• Cross at positive end, arrowhead at negative end

Main Idea

Dipole moment is a measure of the polarity of a molecule. It arises when there is a separation of charge in a molecule due to differences in electronegativity between bonded atoms. The greater the electronegativity difference, the more polar the bond, and the greater the dipole moment.

Core Concept

A dipole is a pair of equal and opposite electric charges separated by a distance. In a polar molecule, one end has a partial positive charge (δ+) and the other end has a partial negative charge (δ−). The dipole moment is a vector quantity — it has both magnitude and direction.

Detailed Explanation

What is Dipole Moment?

• It is the product of the magnitude of positive or negative charges and the distance between them.
• Mathematically, it can be written as:

μ = q × d

• μ = dipole moment
• q = magnitude of charge
• d = distance between the charges

Units of Dipole Moment

• SI unit: Debye (D)
• 1 Debye = 3.335 × 10^-30 C·m (coulomb meter)
• C·m is coulomb meter.

How Electronegativity Difference Determines Dipole Moment

• The difference between the electronegativity of two atoms in a compound determines the overall dipole moment and overall polarity of the compound.
• Greater the electronegativity difference between two atoms, more will be the bond polar (more ionic nature).
• A small difference in electronegativity of two bonded atoms will make the bond less polar or more covalent in nature.

Bond Polarity and Molecular Polarity

• Bond polarity: Determined by the electronegativity difference between two bonded atoms.
• Molecular polarity: Determined by the shape of the molecule and the vector sum of all bond dipoles.

• HCl: H (2.20) and Cl (3.16) → difference = 0.96 → polar bond → polar molecule.
• CO₂: C=O bonds are polar, but the molecule is linear and the dipoles cancel out → nonpolar molecule.

Representation of Dipole Moment

• Dipole moment is represented by an arrow pointing from the positive end to the negative end of the molecule.
• The arrow has a cross at the positive end (δ+) and an arrowhead at the negative end (δ−).

• H is δ+ (positive end)
• Cl is δ− (negative end)
• Arrow points from H to Cl: H → Cl

Factors Affecting Dipole Moment

• Electronegativity difference: Greater difference → greater dipole moment.
• Distance between charges: Greater distance → greater dipole moment.
• Molecular shape: Symmetrical molecules may have zero dipole moment even if bonds are polar.
• Lone pairs: Lone pairs contribute to the dipole moment.

Important Points

• 1. Dipole moment = product of charge and distance.
• 2. μ = q × d
• 3. SI unit: Debye (D)
• 4. Greater electronegativity difference → greater dipole moment.
• 5. Molecular shape determines whether bond dipoles cancel.
• 6. Symmetrical molecules may have zero dipole moment.
• 7. Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar).

Remember

• - Dipole moment: μ = q × d
• - Unit: Debye (D)
• - Greater EN difference: Greater dipole moment
• - Symmetry: Cancels bond dipoles
• - 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).
• 2. Forgetting the vector nature of dipole moment: Dipole moments add as vectors.
• 3. Confusing bond polarity with molecular polarity: Bond polarity is about individual bonds; molecular polarity is about the whole molecule.
• 4. Ignoring lone pairs: Lone pairs contribute to the dipole moment.

Quick Revision

• Dipole moment = product of charge and distance.
• μ = q × d
• SI unit: Debye (D)
• Greater electronegativity difference → greater dipole moment.
• Molecular shape determines whether bond dipoles cancel.
• Symmetrical molecules may have zero dipole moment.
• Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar).

60-Second Revision

• Dipole moment: μ = q × d
• Unit: Debye (D)
• Greater EN difference: Greater dipole moment
• Symmetry: Cancels bond dipoles
• Examples: HCl (polar), CO₂ (nonpolar), H₂O (polar), CCl₄ (nonpolar)

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