CO₂ has two highly polar C=O bonds, yet its measured dipole moment is zero. What best explains this?
ACarbon and oxygen have nearly the same electronegativity, so the C=O bond dipoles are negligible
BThe linear geometry causes the two bond dipole vectors to point in exactly opposite directions and cancel
CAn even number of identical bonds always cancels, regardless of geometry
DCO₂ is an ionic compound, so the concept of bond dipoles does not apply
Each C=O bond is highly polar (O is much more electronegative than C), so both bond dipoles are large. The reason they cancel is geometry: CO₂ is linear, so the two dipole vectors point in exactly opposite directions and sum to zero. If CO₂ were bent like H₂O, the dipoles would add constructively and CO₂ would be polar. This illustrates the core principle: molecular polarity depends on the vector sum of bond dipoles, which is determined by geometry, not just bond polarity.
Question 2 Multiple Choice
A student argues that CCl₄ must be polar because Cl is much more electronegative than C, making each C–Cl bond highly polar. What is wrong with this reasoning?
ANothing — CCl₄ is indeed polar because all four bonds are polar
BThe student is correct about the bonds but ignores that the tetrahedral geometry causes all four bond dipole vectors to cancel, giving a net dipole of zero
CThe student is wrong because C–Cl bonds are actually nonpolar
DThe student confuses ionic character with polarity
The student correctly identifies that each C–Cl bond is polar. The error is ignoring geometry. In tetrahedral CCl₄, the four identical C–Cl bond dipoles point symmetrically outward from the central carbon; their vector sum is exactly zero. Replace one Cl with H (giving CHCl₃) and the symmetry breaks — now the dipoles no longer cancel and the molecule is polar. This is the classic illustration that polar bonds do not guarantee a polar molecule.
Question 3 True / False
A molecule can have polar bonds and still have a net dipole moment of zero.
TTrue
FFalse
Answer: True
This is true, and CO₂ and CCl₄ are the canonical examples. CO₂ has two very polar C=O bonds that cancel because of linear geometry; CCl₄ has four polar C–Cl bonds that cancel because of tetrahedral symmetry. Polarity requires both polar bonds AND an asymmetric arrangement of those bonds such that the vectors do not cancel.
Question 4 True / False
The molecule with the largest individual bond dipoles will typically have the largest molecular dipole moment.
TTrue
FFalse
Answer: False
This is false. Molecular dipole moment is the vector sum of all bond dipoles. A molecule with very large bond dipoles arranged symmetrically (like CCl₄, which has a dipole of 0 D) can have a smaller net dipole than a molecule with smaller but asymmetrically arranged bond dipoles. Geometry — not bond dipole magnitude alone — determines the net result.
Question 5 Short Answer
Why does water (H₂O) have a large net dipole moment while carbon dioxide (CO₂) has zero, even though both molecules contain polar bonds?
Think about your answer, then reveal below.
Model answer: The difference is molecular geometry. Water has a bent geometry (~104.5°), so its two O–H bond dipoles point in directions that partially reinforce each other; their vector sum gives a net molecular dipole of 1.85 D. CO₂ is linear, so its two C=O bond dipoles point in exactly opposite directions and cancel perfectly to zero. The lone pairs on oxygen in water also contribute a dipole component that reinforces the bond dipoles. Same principle applies generally: symmetric geometries cancel bond dipoles; asymmetric geometries produce a net dipole.
The key is vector addition of bond dipoles. In H₂O, the 104.5° angle means the two O–H dipoles are not antiparallel, so they do not cancel. In CO₂, the 180° linear geometry means they are exactly antiparallel and do cancel. Lone pairs add an additional contribution in H₂O, further increasing the dipole. Geometry is the deciding factor.