Question
How do you predict whether a molecule has a net dipole moment? Why does CO₂ have zero dipole moment despite having polar bonds? Why does water have a significant dipole moment? How does molecular geometry determine polarity?
(JEE Main + CBSE Board — reasoning + application)
Solution — Step by Step
Dipole moment () measures the polarity of a molecule. It equals the product of charge separation and bond distance:
Units: Debye (D). C·m.
A molecule has a net dipole moment ONLY if the individual bond dipoles do NOT cancel each other out. This depends entirely on molecular geometry.
| Molecule | Geometry | Bond Polarity | Net Dipole | Reason |
|---|---|---|---|---|
| CO₂ | Linear (180°) | C=O is polar | Zero | Two equal dipoles point in opposite directions → cancel |
| H₂O | Bent (104.5°) | O-H is polar | 1.85 D | Dipoles do NOT cancel due to bent shape |
| BF₃ | Trigonal planar (120°) | B-F is polar | Zero | Three equal dipoles at 120° → cancel |
| NH₃ | Trigonal pyramidal | N-H is polar | 1.47 D | Three dipoles + lone pair → do not cancel |
| CH₄ | Tetrahedral | C-H is slightly polar | Zero | Four equal dipoles in tetrahedral arrangement → cancel |
| CHCl₃ | Tetrahedral (distorted) | C-H and C-Cl differ | Non-zero | Unequal bonds → dipoles do not cancel |
- Draw the Lewis structure
- Determine the molecular geometry (using VSEPR)
- Identify if all surrounding atoms are identical
- If geometry is symmetric AND all atoms are identical → dipole moment = 0
- If geometry is asymmetric OR atoms are different → dipole moment is non-zero
- Lone pairs on the central atom always contribute to the dipole
Why does NH₃ have a dipole moment but BF₃ does not? Both have three bonds to the central atom. The difference: NH₃ has a lone pair on nitrogen that creates an asymmetric electron distribution. BF₃ has no lone pair — its trigonal planar shape is perfectly symmetric.
Lone pairs are like “invisible bonds” that push the electron density in one direction, destroying symmetry.
graph TD
A["Is the molecule symmetric?"] --> B{"All identical atoms around centre?"}
B -->|Yes| C{"Symmetric geometry?"}
C -->|"Linear, Trigonal planar, Tetrahedral, Octahedral"| D["Dipole = 0"]
C -->|"Bent, Pyramidal, See-saw"| E["Dipole ≠ 0"]
B -->|No| F["Dipole ≠ 0"]
G["Lone pair on central atom?"] --> H["Likely asymmetric → Dipole ≠ 0"]
style A fill:#fbbf24,stroke:#000,stroke-width:2px
style D fill:#86efac,stroke:#000
style E fill:#fca5a5,stroke:#000Why This Works
Dipole moment is a vector quantity — it has both magnitude and direction. Individual bond dipoles are vectors pointing from the less electronegative atom toward the more electronegative one. The net dipole is the vector sum of all bond dipoles. In symmetric molecules, these vectors cancel perfectly — like equal forces pulling in opposite directions.
This is why geometry is everything. The same polar bonds (C=O) give zero dipole in CO₂ (linear) but non-zero in carbonyl compounds (bent around carbon in aldehydes/ketones).
Common Mistake
The biggest trap: assuming “polar bonds = polar molecule.” CO₂ has polar C=O bonds but is a non-polar molecule (zero dipole). BF₃ has polar B-F bonds but zero dipole. Polarity of the molecule depends on GEOMETRY, not just bond polarity. Always check whether the bond dipoles cancel before answering. JEE tests this concept repeatedly with molecules like CCl₄ (zero dipole despite polar C-Cl bonds).
Quick check for JEE: any molecule of the form AX_n with no lone pairs on A and all X atoms identical will have zero dipole moment if the geometry is linear (n=2), trigonal planar (n=3), tetrahedral (n=4), or octahedral (n=6). The moment lone pairs appear or X atoms differ, the dipole becomes non-zero.