Question

The Lewis structure of \(\mathrm{N}_{2}\) indicates that the nitrogento-nitrogen bond is a triple covalent bond. Other evidence suggests that the \(\sigma\) bond in this molecule involves the overlap of \(s p\) hybrid orbitals. (a) Draw orbital diagrams for the N atoms to describe bonding in \(\mathrm{N}_{2}\) (b) Can this bonding be described by either \(s p^{2}\) or \(s p^{3}\) hybridization of the \(\mathrm{N}\) atoms? Can bonding in \(\mathrm{N}_{2}\) be described in terms of unhybridized orbitals? Explain.

Short answer

The triple bond in N2 is formed by one sigma bond and two pi bonds which are the result of sp hybridization of N atoms. The bonding in N2 cannot be accurately described using only unhybridized orbitals.

Step-by-step solution

Draw the Lewis Structure for N2

The nitrogen atom has 5 valence electrons. As there are two nitrogen atoms, we have a total of 10 electrons for N2. We'll put two electrons in the nitrogen-nitrogen bond forming a single covalent bond and then place pairs of electrons around each nitrogen atom. After fulfilling the octet for each nitrogen atom, we still have 2 more electrons, which we will place in the bond, forming a triple covalent bond. The resulting Lewis structure of N2 shows a triple bond between two nitrogen atoms.

Draw Orbital Diagrams for N atoms

Each Nitrogen atom (N) has two 1s, two 2s, and three 2p electrons. The 1s and 2s are core electrons and do not participate in bonding. The 2p electrons do. Nitrogen's 2s and three 2p orbitals hybridize (combine) to form four identical orbitals. These orbitals align themselves in space in such a way as to minimize repulsion between their electron clouds. Three of these orbitals form sigma bonds with a hydrogen atom's 1s orbital. The fourth hybrid orbital holds the remaining pair of nonbonding electrons.

Discuss the Type of Hybridization in N2

In N2, the bonding cannot be described by either sp2 or sp3 hybridization. Instead, it can be described by sp hybridization. This is because when a bond is formed, the s orbital and one p orbital of each nitrogen atom hybridize to form two sp hybrid orbitals. These sp orbitals overlap to form the sigma bond in N2. The remaining p orbitals on each nitrogen form two pi bonds. Therefore, the triple bond in N2 is composed of one sigma bond and two pi bonds, and is the result of sp hybridization. Also, it's important to note that the bonding in N2 cannot be accurately described using only unhybridized orbitals since the bonding primarily involves the sp hybrid orbitals.

Key concepts

Lewis Structure

A Lewis structure is a diagram that represents the arrangement of electrons in a molecule. It helps us visualize how atoms in a molecule are bonded together and how electrons are distributed among them. In the case of \(_2\), each nitrogen atom has 5 valence electrons. By drawing the Lewis structure for \(_2\), you can illustrate a triple bond between the two nitrogen atoms.

• Start by connecting the atoms with a single bond.
• Then, add electrons to satisfy the octet rule, ensuring each nitrogen has 8 electrons.
• The pair of nitrogen atoms share three pairs of electrons, forming a triple bond, resulting in the most stable arrangement.

This representation is crucial for understanding molecular geometry and reactivity.

Hybridization

Hybridization is a concept in chemistry where atomic orbitals mix to form new hybrid orbitals. These are used to describe the shape of electron clouds around atoms that are bonded. For \(_2\), hybridization helps us explain the strength and length of the bonds.

• In \(_2\), sp hybridization occurs. This involves the mixing of one s orbital and one p orbital to form two equivalent sp orbitals.
• Each nitrogen atom in \(_2\) uses one of these sp orbitals to form a sigma bond with the other.

This hybridization results in a linear arrangement, which is key due to the strength of the triple bond present.

Triple Bond

Triple bonds are a type of covalent bond involving the sharing of three pairs of electrons between two atoms. In \(_2\), the triple bond between the nitrogen atoms provides a good example of how strong covalent bonds can be.

• A triple bond consists of one sigma bond and two pi bonds.
• Sigma bonds are formed by the head-on overlap of orbitals, while pi bonds result from the side-to-side overlap.

The presence of the triple bond in \(_2\) makes it incredibly stable and difficult to break, reflecting nitrogen's relatively inert nature under normal conditions.

Sigma and Pi Bonds

Sigma (\(\sigma\)) and pi (\(\pi\)) bonds are two types of covalent bonds. Understanding these bonds is crucial for grasping molecular orbital theory and stability.

Sigma Bonds

Sigma bonds form from the head-on overlap of orbitals, such as sp hybrid orbitals in \(_2\).

• They are the strongest type of covalent bonds due to maximum overlap.
• Sigma bonds are always involved in single bonds.

Pi Bonds

Pi bonds arise from the side-to-side overlap of parallel p orbitals.

• Each pi bond is weaker than a sigma bond; hence two pi bonds accompany the sigma bond in a triple bond.
• They add to the bond energy and stability of the molecule, yet allow for some flexibility in shape.

In \(_2\), you have one sigma bond and two pi bonds, providing a comprehensive understanding of its strong and stable triple bond nature.