Question

Draw the lewis structure for \(\mathrm{NF}_{3}\). What are its electron-pair and molecular geometries? What is the hybridization of the nitrogen atom? What orbitals on \(\mathrm{N}\) and \(\mathrm{F}\) overlap to form bonds between these elements?

Short answer

NF3 has a trigonal pyramidal molecular geometry and an sp3 hybridization of nitrogen. N-F bonds involve overlap between nitrogen's sp3 hybrid orbitals and fluorine's p orbitals.

Step-by-step solution

Count Valence Electrons

Determine the total number of valence electrons in \(\mathrm{NF}_3\). Nitrogen (N) has 5 valence electrons, and each fluorine (F) has 7 valence electrons. Therefore, the total is \(5 + 3 \times 7 = 26\) valence electrons.

Arrange Atoms and Assign Electrons

Place the nitrogen atom in the center and arrange three fluorine atoms around it. Connect each fluorine to nitrogen with a single bond, using 2 electrons per bond. This uses 6 electrons, leaving 20 electrons to be distributed.

Complete Octets for Fluorine Atoms

Each fluorine atom needs 6 more electrons to complete its octet. Distribute these electrons around the three fluorine atoms (6 electrons per F), using 18 electrons total. Now, all fluorines have a complete octet.

Place Remaining Electrons on Nitrogen

There are 2 remaining electrons after completing the fluorine octets, which will be placed as a lone pair on the nitrogen atom.

Determine Electron-Pair Geometry

Including the lone pair, nitrogen is surrounded by four areas of electron density, resulting in a tetrahedral electron-pair geometry.

Determine Molecular Geometry

With one lone pair and three bonded pairs, the molecular geometry of \(\mathrm{NF}_3\) is trigonal pyramidal.

Determine Hybridization of Nitrogen

To accommodate four regions of electron density (three bonds and one lone pair), the nitrogen undergoes \(sp^3\) hybridization.

Identify Overlapping Orbitals

The \(sp^3\) hybrid orbitals of nitrogen overlap with the \(2p\) orbitals of each fluorine to form the \(N-F\) sigma bonds.

Key concepts

Electron-Pair Geometry

The electron-pair geometry of any molecule is determined by the spatial arrangement of regions of electron density (bonding and non-bonding) around the central atom. In the case of nitrogen trifluoride (\(\mathrm{NF}_{3}\)), the focus is on the nitrogen atom. This central atom is surrounded by three fluorine atoms forming bonded pairs and one lone pair. A total of four regions of electron density indicates a tetrahedral electron-pair geometry. This arrangement helps minimize electron pair repulsions, which is a key principle of Valence Shell Electron Pair Repulsion (VSEPR) theory.

• Tetrahedral Arrangement: When looking at the electron-pair geometry, it's all about considering every electron region, counting both bonds and lone pairs.
• Understanding Repulsion: The tetrahedral shape spreads out these regions evenly to minimize interaction.

Molecular Geometry

While electron-pair geometry considers all electron densities, molecular geometry describes the shape formed only by the atoms. For \(\mathrm{NF}_{3}\), molecular geometry is influenced by the one lone pair of electrons on nitrogen. This lone pair occupies more space than the bonded pairs, pushing the three fluorine atoms closer together. As a result, the molecular geometry of \(\mathrm{NF}_{3}\) is trigonal pyramidal.

• Lone Pair Influence: Lone pairs exert more repulsion than bonded pairs, thus altering bond angles.
• Visualizing the Shape: The three fluorine atoms and the lone pair create a pyramid with nitrogen at the peak.

Hybridization

Hybridization is the process by which atomic orbitals mix to form new hybrid orbitals, which are used to form bonds. In the case of \(\mathrm{NF}_{3}\), the nitrogen atom undergoes \(\text{sp}^3\) hybridization. This occurs when one s orbital and three p orbitals in nitrogen mix to create four equivalent hybrid orbitals.

• Identifying Hybridization: Considers both bonded pairs and lone pairs, totaling four regions of electron density.
• Creating Hybrid Orbitals: These hybrid \(\text{sp}^3\) orbitals are evenly oriented in space, accommodating both bonding and non-bonding interactions.

Overlap of Orbitals

The bonds between nitrogen and fluorine in \(\mathrm{NF}_{3}\) are formed through the overlap of their orbitals. Each nitrogen \(\text{sp}^3\) hybrid orbital overlaps with a fluorine \(2p\) orbital. This type of overlap leads to the creation of sigma bonds, which are the strongest covalent bonds and involve head-on overlap.

• Sigma Bond Formation: This results from the direct overlap of orbitals, it's strong due to its high electron density along the bond axis.
• Role in Bond Strength: The effective overlap ensures stable and robust \(N-F\) bonds.