Why does the phosphorus trifluoride molecule have dipole moment? My book says that from drawing the Lewis structure for the molecule below, you can conclude that it has dipole moment. Can someone please explain why?



I understand that the dipole moment expresses the polarity of the bond. And that the dipole moment of a molecule is the sum of the dipole moments for each bond in the molecule.
Fluorine is more electronegative than phosphorus.
The phosphorus is at the apex of a pyramid, the base of the pyramid being an equilateral triangle with a fluorine atom at each vertex.
The F-P-F angles are 96 degrees.
Each F-P bond contributes to the net dipole moment, as a vector from the P to the F.  The net dipole moment is the vector sum of the vectors along the three P-F bonds.
If all four atoms were in a plane (a trigonal planar geometry), there would be no net dipole moment.  But because the geometry is pyramidal, the vector sum is not zero and there is a net dipole moment.
In some molecules lone pairs also contribute to the net dipole moment, in this particular case the lone pair of phosphorus is in a mostly s-like orbital, so in first approximation it can be neglected.