Boron trifluoride (BF3) is a nonpolar molecule. While it can be confusing because the individual bonds inside the molecule are highly polar, the overall molecule ends up with a net dipole moment of zero. Here is a breakdown of why this happens:
1. Highly Polar Bonds
Fluorine is the most electronegative element on the periodic table, whereas boron has a much lower electronegativity. Because of this massive difference, the electrons in each $\text{B-F}$ bond are pulled strongly toward the fluorine atoms, making each individual bond highly polar.
2. Symmetrical Geometry
According to the VSEPR (Valence Shell Electron Pair Repulsion) theory, boron has three valence electrons and forms three single bonds with no remaining lone pairs. This forces the molecule into a trigonal planar geometry, where all three fluorine atoms are spaced perfectly 120° apart in a single flat plane.
3. Dipole Cancellation
Because the trigonal planar shape is perfectly symmetrical, the outward pulling forces (dipole vectors) of the three $\text{B-F}$ bonds cancel each other out completely. Think of it like a three-way tug-of-war where everyone is pulling with the exact same strength at equal angles—nobody moves anywhere.