Questions: Polar Covalent Bonds and Dipole Moments

Question 1

Bond A connects two atoms with an electronegativity difference of 1.5 and a bond length of 100 pm. Bond B connects two atoms with an electronegativity difference of 0.8 and a bond length of 200 pm. Which bond necessarily has the larger dipole moment?

• A. Bond A, because a larger electronegativity difference always produces a larger dipole moment
• B. Bond B, because a longer bond length means the charges are farther apart, increasing the dipole moment
• C. Cannot be determined without calculating μ = q × d for both bonds using actual partial charge values
• D. They are equal because the effects of electronegativity difference and bond length cancel out

Answer: (C)

Dipole moment is μ = q × d: the product of charge separation (proportional to, but not exactly equal to, electronegativity difference) and distance. Both factors matter independently. Bond A has larger charge separation but shorter distance; Bond B has smaller charge separation but longer distance. Without the actual partial charge magnitudes, we cannot conclude which product is larger. This tests whether students understand that dipole moment is NOT determined by electronegativity difference alone — both q and d must be considered.

Question 2

In a polar covalent bond between atoms X and Y (where Y is more electronegative), which statement correctly describes the conventional dipole moment vector?

• A. The arrow points from Y toward X, because the partial positive charge is located on X
• B. The arrow points from X toward Y, because convention places the arrow from δ+ to δ−
• C. The arrow points from Y toward X, showing the direction electrons flow during bond formation
• D. The vector has no defined direction because dipole moment is a scalar quantity

Answer: (B)

By convention, the dipole moment arrow points from the positive end (δ+) to the negative end (δ−). Since Y is more electronegative, electron density shifts toward Y, making Y the δ− end and X the δ+ end. The arrow therefore points from X toward Y — from δ+ to δ−. Option A reverses the direction. Option C confuses the direction of electron flow with the conventional dipole vector. Option D is wrong — dipole moment is a vector, not a scalar, with both magnitude and direction.

Question 3

A bond between two identical atoms (such as Cl–Cl) has a dipole moment of exactly zero.

• T. True
• F. False

Answer: True

When two identical atoms share a bond, there is no electronegativity difference, so electron density is perfectly symmetric between them. No partial charges (δ+ or δ−) develop on either atom. The dipole moment formula μ = q × d gives zero because q = 0. This is the defining case of a pure nonpolar covalent bond.

Question 4

The dipole moment of a bond depends mainly on the electronegativity difference between the bonded atoms, not on the bond length.

• T. True
• F. False

Answer: False

Dipole moment is the product of charge separation AND distance: μ = q × d. Bond length (d) is an independent factor. A bond with a modest electronegativity difference but a very long bond length can have a larger dipole moment than a bond with a large electronegativity difference but a very short bond length. Both the magnitude of the partial charges and the physical distance between them must be considered.

Question 5

Why is the dipole moment described as a vector quantity, and why does this matter when predicting whether a molecule as a whole is polar?

Model answer: A vector has both magnitude and direction, unlike a scalar which has only magnitude. Each bond contributes a dipole vector pointing from δ+ to δ−. To find the overall molecular dipole moment, these individual bond dipole vectors must be added as vectors — they may reinforce or cancel depending on the molecular geometry. CO₂ has two large, polar C=O bond dipoles that point in exactly opposite directions and cancel completely, making the molecule nonpolar overall despite having two polar bonds. H₂O has two O-H bond dipoles that point in directions that partially reinforce each other, producing a net molecular dipole.

The vector nature of dipole moment is what connects bond-level polarity to molecular-level polarity. Students who treat dipole moment as a scalar (just a number) will incorrectly predict that any molecule with polar bonds must be polar overall, missing the key role of geometry in determining whether bond dipoles cancel or add.