I had to clear this up since there were many threads with the wrong analysis of this problem. Lone pair electrons DO have and affect on dipole moments. Lone pair electrons do have an effect on the net dipole moment of a molecule. A dipole moment is the product of the magnitude of the partial charges on the molecule and the distance in which they are separated. Lone pair electrons affect both the partial charges and the distance between partial charges. The NH3 and NF3 are perfect examples yet are explained wrong here. Using electronegativities NF3 should have a greater dipole moment because N=3.04, F=3.98, and H=2.20. Thus the greater disparity is between N and F. However, the dipole moment of NF3 is quite small at 0.23 D (D is debye). The dipole moment of NH3 is 1.47 D. The lone pair of electrons extends into space and increases the charge separation in the molecule, thus increasing its dipole (see definition above). You can see electronegativities do NOT consider everything that must be considered for dipole moments. Lone pairs DO have an effect on the net dipole moment of a molecule (a.k.a. the polarity of the molecule). However, for most applications in a classroom lone pair effects on a dipole moments are not considered for sake of simplicity.