The molecular geometry of nitrogen trifluoride (NF3) can be described using VSEPR (Valence Shell Electron Pair Repulsion) theory.
Nitrogen (N) is at the center, surrounded by three fluorine (F) atoms. Nitrogen has five valence electrons, and when it forms three bonds with fluorine, it uses three of its electrons and results in one lone pair remaining on the nitrogen atom.
Since we have a lone pair on the nitrogen, the molecular geometry is not tetrahedral (which would be the case with no lone pairs) but rather trigonal pyramidal. In this arrangement, the lone pair occupies more space than the bond pairs, causing a distortion in the ideal bond angles.
Theoretical bond angles in a perfect tetrahedral geometry are 109.5°. However, due to the presence of the lone pair on the nitrogen, the bond angles in NF3 are slightly less than 109.5°, typically about 102° to 103° between the fluorine atoms.
In summary, NF3 features trigonal pyramidal molecular geometry with bond angles of approximately 102° to 103° due to the influence of the lone pair on the nitrogen atom. This structure not only affects the angles but also influences the overall polarity and reactivity of the molecule.