A Lewis structure is a simple diagram used to represent the arrangement of electrons in a molecule. This diagram illustrates how atoms bond together, showing both shared pairs (chemical bonds) and unshared electrons (lone pairs). By depicting the valence electrons, Lewis structures help predict a molecule’s properties and its potential to interact with other substances. The molecule of interest is Nitrogen Trifluoride, which has the chemical formula \(\text{NF}_3\).
Calculating the Necessary Components
The first step in determining the Lewis structure is calculating the total number of valence electrons available. Nitrogen (N) belongs to Group 15, contributing five valence electrons. Fluorine (F) belongs to Group 17, and the three fluorine atoms contribute seven valence electrons each. This yields a total of \(5 + (3 \times 7) = 26\) valence electrons for the entire \(\text{NF}_3\) molecule.
Next, the central atom must be identified, which is the atom to which all others will bond. The least electronegative atom typically occupies this central position. Since fluorine is the most electronegative element, nitrogen is the central atom in \(\text{NF}_3\). The three fluorine atoms are then arranged symmetrically around the central nitrogen atom.
Drawing the Lewis Structure for \(\text{NF}_3\)
The process begins by connecting the three fluorine atoms to the central nitrogen atom using single covalent bonds. Since each single bond uses two shared electrons, \(3 \times 2 = 6\) electrons are used for the skeletal structure. Subtracting these six bonding electrons from the total of 26 leaves 20 electrons remaining to be distributed.
These remaining electrons are first used to satisfy the octet rule for the terminal fluorine atoms. Each fluorine atom already has two electrons from its bond with nitrogen, requiring six more electrons to achieve a stable octet. Placing three lone pairs (six electrons) on each of the three fluorine atoms consumes \(3 \times 6 = 18\) electrons.
After accounting for the 6 bonding electrons and the 18 electrons on the fluorine atoms, two electrons remain (\(26 – 6 – 18 = 2\)). These final two electrons are placed as a single lone pair on the central nitrogen atom. The finished Lewis structure shows nitrogen bonded to three fluorine atoms, with the nitrogen atom having one lone pair and each fluorine atom having three lone pairs.
Molecular Geometry and Polarity
The completed Lewis structure provides the necessary information to determine the molecule’s three-dimensional shape and its polarity (overall charge distribution). The shape is determined by the Valence Shell Electron Pair Repulsion (VSEPR) theory. VSEPR states that electron groups—both bonding pairs and lone pairs—arrange themselves to be as far apart as possible. The central nitrogen atom has four groups of electrons surrounding it: three single bonds and one lone pair.
This arrangement of four electron groups suggests a tetrahedral electron geometry. However, the molecular geometry, which describes only the position of the atoms, is different because the lone pair occupies space but is not an atom. The repulsive force of the lone pair pushes the three N-F bonds downward. This results in a molecular shape known as trigonal pyramidal, with the nitrogen atom at the apex and the fluorine atoms forming the base.
The \(\text{NF}_3\) molecule is polar, meaning it has a net positive and net negative end. Each N-F bond is individually polar because fluorine is more electronegative than nitrogen, pulling the bonding electrons toward itself. In the trigonal pyramidal shape, the dipole moments of the three N-F bonds point away from the nitrogen atom toward the base of the pyramid. The lone pair on the nitrogen atom creates a separate, strong dipole moment pointing upward. Although the lone pair’s dipole partially counteracts the bond dipoles, the asymmetry of the trigonal pyramidal shape prevents complete cancellation, resulting in a small, net dipole moment.