Question

The structure of \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\) involves a planar arrangement of \(\mathrm{N}\) and \(\mathrm{Si}\) atoms, whereas that of the related compound \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\) has a pyramidal arrangement of N and \(\mathrm{C}\) atoms. Propose bonding schemes for these molecules that are consistent with this observation.

Short answer

Both \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\) and \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\) are sp3 hybridized. The difference in the geometric structures is because of the difference in the size and electronegativity of Silicon and Carbon. Due to the larger size and less electronegativity of Silicon, \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\) has a planar configuration. \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\), on the other hand, has a pyramidal structure due to the smaller size and higher electronegativity of Carbon.

Step-by-step solution

Determine the central atom

Identify the central atom in the molecule. In both compounds \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\) and \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\), Nitrogen (N) acts as the central atom.

Count the valence electrons

Count the valence electrons for each central atom. The Nitrogen atom has 5 valence electrons. Each Hydrogen atom contributes 1 valence electron, and Carbon contributes 4, hence \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\) has total of 8 valence electrons. Silicon in the case of \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\), contributes 4 valence electrons, which gives a total of 8 valence electrons.

Determine hybridization and molecular geometry

The hybridization of an atom in a molecule can be calculated by the formula: \( \text{hybridization} = 1/2 * (\text{Number of valence electrons} + \text{Number of monovalent atoms attached to the atom} - \text{Charge on cation} + \text{Charge on anion})\). Applying this formula, both \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\) and \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\) show sp3 hybridization. However, due to the larger size and less electronegativity of Silicon as compared to Carbon, the bond pair-bond pair repulsion is much lesser, leading to a planar configuration in the case of \(\mathrm{N}\left(\mathrm{SiH}_{3}\right)_{3}\). Whereas in \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3}\), due to higher electronegativity of Carbon, the bond pair-bond pair repulsion is greater, leading to a pyramidal shape.

Key concepts

Hybridization

Hybridization is a key concept in molecular geometry, determining the shape of molecules. It involves the mixing of atomic orbitals to form new hybrid orbitals. These hybrid orbitals have different energies and shapes than the original orbitals and lead to the formation of bonds in molecules.
In the case of both \(\mathrm{N(SiH_3)_3}\) and \(\mathrm{N(CH_3)_3}\), the central nitrogen atom undergoes \(sp^3\) hybridization. This type of hybridization involves one s orbital and three p orbitals mixing to create four equivalent \(sp^3\) hybrid orbitals.

• Each \(sp^3\) hybrid orbital can form a bond with another atom, optimizing the molecule's geometry.
• The geometry resulting from \(sp^3\) hybridization is typically tetrahedral.

However, the ultimate shape of the molecule may vary, as seen in these compounds, due to differences in bond repulsions, which are influenced by the characteristics of the atoms bonded to nitrogen.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and are crucial in determining how atoms interact and bond with each other. In molecules, understanding the valence electrons helps predict the possible structures and properties thru bonding.
For nitrogen in \(\mathrm{N(SiH_3)_3}\) and \(\mathrm{N(CH_3)_3}\), nitrogen has 5 valence electrons. Silicon (\(\mathrm{Si}\)), like carbon (\(\mathrm{C}\)), each have 4 valence electrons.

• Each hydrogen atom contributes an additional valence electron in the bonding process.
• Every \(\mathrm{SiH_3}\) or \(\mathrm{CH_3}\) group adds electrons to the total, driving the bonding arrangement.

Evaluating the total count of valence electrons aids in predicting molecular geometry and helps in determining the need for hybridization to accommodate the electrons efficiently.

Bond Pair Repulsion

The concept of bond pair repulsion is pivotal in understanding molecular shape, especially when predicting whether a molecule will form a planar or a pyramidal structure. This occurs because the electrons in bonds repel each other, influencing how the bonds orient themselves in space.
In \(\mathrm{N(SiH_3)_3}\), the silicon atoms are less electronegative compared to carbon as seen in \(\mathrm{N(CH_3)_3}\). This lower electronegativity in silicon reduces the electronegative pull on the nitrogen atom, thus leading to a smaller bond pair-bond pair repulsion.

• Lower repulsion allows the molecule to adopt a planar configuration, as with \(\mathrm{N(SiH_3)_3}\).
• In contrast, \(\mathrm{N(CH_3)_3}\) experiences greater bond pair repulsion due to the higher electronegativity of carbon, resulting in a pyramidal structure.

Understanding how different atoms contribute to bond repulsion helps explain why molecules with similar hybridizations can have different geometries.