Dipole Moment: MDCAT Chemistry notes
Dipole Moment for MDCAT: polar and non-polar bonds, electronegativity difference, molecular symmetry, zero-dipole molecules and dipole-dipole forces.
Polar and non-polar bonds
When two atoms share electrons equally, the bond is non-polar. This happens between identical atoms, as in homonuclear diatomic molecules such as $\mathrm{H_2}$, $\mathrm{Cl_2}$ and $\mathrm{N_2}$. When one atom is more electronegative, it pulls the shared pair towards itself. The bond is then polar, with partial charges $\delta^+$ and $\delta^-$.
- The electronegativity difference between the bonded atoms is an index of the polar nature of the bond. The larger the difference, the more polar (more ionic) the bond.
- C–O in methanol is a polar covalent bond, because O is much more electronegative than C.
Dipole moment
The dipole moment measures polarity. It is a vector quantity, pointing from the positive end to the negative end in the chemistry convention.
Key formulas
$$\mu = q \times d$$Here $q$ is the partial charge and $d$ the distance between the charges. The unit is the debye (D), where $1\ \mathrm{D} = 3.336 \times 10^{-30}$ C m. The SI unit is the coulomb metre.
Molecular polarity: symmetry decides
A molecule's dipole moment is the vector sum of its bond dipoles. A molecule with polar bonds can still have zero dipole moment if its shape is symmetrical, because the bond dipoles cancel.
| Zero dipole moment (non-polar molecule) | Non-zero dipole moment (polar molecule) |
|---|---|
| $\mathrm{CO_2}$, $\mathrm{BeCl_2}$ (linear) | HCl (single polar bond) |
| $\mathrm{BF_3}$, $\mathrm{AlCl_3}$ (trigonal planar) | $\mathrm{H_2O}$ (bent) |
| $\mathrm{CH_4}$, $\mathrm{CCl_4}$ (regular tetrahedral) | $\mathrm{NH_3}$, $\mathrm{NF_3}$ (pyramidal) |
| $\mathrm{H_2}$, $\mathrm{Cl_2}$ (identical atoms) | $\mathrm{CH_3Cl}$, $\mathrm{CH_3OH}$ (unsymmetrical) |
Uses of dipole moment: it shows whether a molecule is polar, and it gives information about its shape. $\mathrm{CO_2}$ has zero dipole moment, so it must be linear. $\mathrm{H_2O}$ has a dipole moment of about 1.85 D, so it must be bent.
Ionic or covalent? Polarization
A small, highly charged cation distorts (polarizes) the electron cloud of a nearby anion. This gives the bond partial covalent character. $\mathrm{Al^{3+}}$ is small and triply charged, so $\mathrm{AlCl_3}$ is largely covalent. NaCl, KCl and $\mathrm{MgCl_2}$ are ionic.
Dipole–dipole forces
Polar molecules attract one another through dipole–dipole forces: the $\delta^+$ end of one molecule attracts the $\delta^-$ end of the next. Ammonia (polar) shows them. Methane and ethane (non-polar) and neon (a single atom) have only London dispersion forces.
Common MDCAT traps
- Polar bonds do not always mean a polar molecule. Symmetry can cancel them, as in $\mathrm{CO_2}$, $\mathrm{BF_3}$ and $\mathrm{CCl_4}$.
- $\mathrm{NF_3}$ and $\mathrm{NH_3}$ are pyramidal, so they are polar, unlike planar $\mathrm{BF_3}$.
- Strictly, a non-polar bond exists only between identical atoms. C=O bonds in $\mathrm{CO_2}$ are polar even though the molecule is non-polar. Read carefully whether the question asks about the bond or the molecule.
- $\mathrm{AlCl_3}$ is covalent despite containing a metal.
Quick revision
- $\mu = q \times d$, measured in debye.
- Symmetrical molecules have $\mu = 0$.
- A larger electronegativity difference means a more polar bond.
- Homonuclear diatomic molecules have non-polar bonds.
- Polar molecules show dipole–dipole attractions.