Identifying Polar Molecules
A polar molecule is one with a permanent Dipole Moment. A polar molecule must have a slightly
positive end opposite a slightly negative one. This cannot happen if the molecule is too symmetric! If
a molecule is 'spherical' enough, then each end of the molecule will have the same properties and in
must be non-polar.
Examples of molecules that have exactly zero dipole moment and therefore be referred to as nonpolar by symmetry are:
Homonuclear Diatomics, and molecules with a center of inversion, eg. H2 and CO2
Symmetric 'Disk" shaped molecules (Trigonal Planar, Square Planar, Hexagonal Planar, etc), eg.
Benzene (C6H6) and BCl3
Symmetric 'Ball' shaped molecules (Tetrahedral and Octahedral), eg. Methane (CH4) and SF6
Note: These arguments only hold for symmetrically substituted molecules; Asymmetric substitution
giverise to a net dipole.
Molecules that have 'low' symmetry will always have at least a small dipole moment and therefore be
referred to as polar. Examples of such low symmetry molecular shapes include:
Bent molecules, eg. water
T-shaped molecules, eg. IF3
See-Saw molecules, eg. SF4
Pyramidal molecules (trigonal pyramidal, square pyramidal, etc) [NOT bi-pyramidal], eg. NF3,
BrF5
When molecules are less than 'perfectly' symmetric, a dipole moment results from the unequal
sharing of the electrons between bonded atoms. Remember what indicated how unequal the sharing
of electrons was? Yup, the electronegativity:
http://www.chem.ufl.edu/~itl/2045/lectures/lec_16.html
We normally refer to the charge separation resulting from the unequal sharing of the electrons in a
chemical bond as a Bond Dipole. The greater the difference in Electronegativities between the
bonded atoms, the greater the Bond Dipole.
The molecule's overall dipole moment is the result of the vector sum of all the bond dipoles within it.
In the symmetric molecules above, no matter what the bond dipoles are, the net dipole moment of the
molecule is zero because the bond dipoles cancel.
Lone pairs contribute to the molecule's dipole moment even though the do not constitute a 'bond'.
Clearly the nucleus 'end' of the lone pair is positive and the electron 'end' is negative so one might
think of a 'lone pair dipole' contributing to the polarity of the molecule in analogy to a bond dipole.
This behavior is demonstrated in the relative magnitudes ad directions of the dipole moments in the
molecules PH3 (=0.58 D), NH3 (=1.47 D), and PF3 (=1.03 D)
http://www.chem.ufl.edu/~itl/2045/lectures/lec_16.html