Dipole moments describe how electric charge is unevenly distributed in a chemical bond or an entire molecule. They matter because polarity helps explain boiling points, solubility, intermolecular forces, and how molecules interact in biology and materials. A bond becomes polar when two bonded atoms attract shared electrons unequally.
The atom with greater electronegativity becomes partially negative, while the other becomes partially positive.
A molecular dipole is the vector sum of all bond dipoles and any effects from lone pairs. Even if a molecule has polar bonds, its overall dipole can cancel if the molecule is highly symmetrical, such as CO2 or CCl4. In contrast, bent, trigonal pyramidal, or otherwise asymmetrical shapes often produce a net dipole, such as in H2O and NH3.
Predicting polarity requires combining electronegativity differences with molecular geometry rather than looking at bonds alone.
Key Facts
- A bond dipole points toward the more electronegative atom.
- Dipole moment is often written as μ = q × r, where q is separated charge and r is distance.
- Partial charges are labeled δ+ on the less electronegative atom and δ− on the more electronegative atom.
- A molecule is polar if its bond dipoles and lone pair effects add to a nonzero net dipole.
- Symmetrical molecules can be nonpolar even when they contain polar bonds, such as CO2 and CCl4.
- Electronegativity difference helps estimate bond polarity: larger ΔEN usually means a more polar bond.
Vocabulary
- Bond dipole
- A bond dipole is the separation of partial positive and partial negative charge across a covalent bond.
- Molecular dipole moment
- A molecular dipole moment is the overall polarity of a molecule found by adding all bond dipoles as vectors.
- Electronegativity
- Electronegativity is an atom's ability to attract shared electrons in a chemical bond.
- Partial charge
- A partial charge is a small charge imbalance shown with δ+ or δ− rather than a full ionic charge.
- Molecular geometry
- Molecular geometry is the three-dimensional arrangement of atoms in a molecule.
Common Mistakes to Avoid
- Assuming every molecule with polar bonds is polar, which is wrong because symmetrical shapes can make bond dipoles cancel.
- Drawing the dipole arrow toward δ+, which is wrong in chemistry because the arrow points toward the more electronegative δ− atom.
- Ignoring lone pairs when predicting polarity, which is wrong because lone pairs affect molecular shape and can create an uneven charge distribution.
- Using electronegativity difference alone to decide molecular polarity, which is wrong because the three-dimensional geometry determines whether dipoles add or cancel.
Practice Questions
- 1 HCl has a bond length of 1.27 × 10^-10 m and an effective separated charge of 3.33 × 10^-20 C. Calculate the bond dipole moment using μ = q × r.
- 2 A molecule has two equal bond dipoles of 1.5 D pointing in exactly opposite directions. What is the net dipole moment?
- 3 CO2 has two polar C=O bonds, while H2O has two polar O-H bonds. Explain why CO2 is nonpolar but H2O is polar.