Question

The molecules \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) are isoelectronic but their properties are quite different. Although as a first approximation we often use the same MO diagram for both, suggest how the \(\mathrm{MOs}\) in \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) might be different.

Short answer

Despite having the same number of electrons, the molecular orbitals of \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) are influenced by the electronegativity differences between their constituent atoms. In \(\mathrm{N}_{2}\), the bond order is 3, and it is quite stable due to its symmetric MO diagram. In \(\mathrm{CO}\), the bond order is also 3, but it has a polar bond due to the difference in electronegativity between carbon and oxygen atoms. The polarity of \(\mathrm{CO}\) results in a more stable \(\sigma_{2p}\) orbital compared to that in \(\mathrm{N}_{2}\), leading to different reactivity and properties despite their isoelectronic nature.

Step-by-step solution

Understand Molecular Orbital Theory

Molecular Orbital (MO) Theory is a method for determining molecular structure and properties by utilizing the wave functions of individual atoms to describe the behaviour of electrons in molecules. It explains how atomic orbitals combine to form molecular orbitals, which can hold a specific number of electrons and, thus, determine the bond order, bond strength, and magnetic properties of the molecule.

Recognize the Isoelectronic Nature

\(\mathrm{N}_{2}\) and \(\mathrm{CO}\) are isoelectronic molecules, which means they have an equal number of electrons. The total number of electrons in \(\mathrm{N}_{2}\) is 14 (7 from each nitrogen atom), while in \(\mathrm{CO}\), there are 6 electrons from carbon and 8 electrons from oxygen, also totaling 14 electrons.

Create MO Diagrams

To compare the MOs of \(\mathrm{N}_{2}\) and \(\mathrm{CO}\), we need to create MO diagrams for both molecules. We know that both molecules have 14 electrons, which fill the MOs in the following order: \(\sigma_{1s}\), \(\sigma_{1s}^*\), \(\sigma_{2s}\), \(\sigma_{2s}^*\), \(\sigma_{2p}\), \(\pi_{2p}\), \(\pi_{2p}^*\), and \(\sigma_{2p}^*\). Note that due to electronegativity differences between elements, the MO diagrams may not be identical between the two.

Compare the Energy Levels of the Orbitals

In the MO diagrams of both \(\mathrm{N}_{2}\) and \(\mathrm{CO}\), the shapes and energy levels of the molecular orbitals are influenced by the electronegativity of the atoms. Nitrogen atoms have similar electronegativity values resulting in a symmetric MO diagram, while the stronger electronegativity of oxygen in \(\mathrm{CO}\) can lead to energy levels of the MO orbitals being less symmetric. For example, in \(\mathrm{CO}\), the \(\sigma_{2p}\) orbital might be more stabilized (lower in energy) than in \(\mathrm{N}_{2}\) because of the greater involvement of more electronegative oxygen's \(2p\) orbital.

Explain the Differences in Properties

Despite having the same number of electrons, the molecular orbitals of \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) are influenced by the electronegativity differences between their constituent atoms. This affects their bond strength, bond length, and bond order. In \(\mathrm{N}_{2}\), the bond order is 3 (the molecule has a triple bond) and it is quite stable due to its entirely symmetric MO diagram. In \(\mathrm{CO}\), the bond order is also 3 (triple bond), but it has a polar bond due to the difference in electronegativity between the carbon and oxygen atoms. The polarity of \(\mathrm{CO}\) also results in a more stable \(\sigma_{2p}\) orbital compared to that in \(\mathrm{N}_{2}\). This leads to different reactivity and properties of the two molecules, despite their isoelectronic nature.

Key concepts

Isoelectronic molecules

Isoelectronic molecules are molecules having the same number of electrons, even though they might contain different types of atoms. This fascinating concept forms an essential part of understanding molecular behavior. For instance, in the case of \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) molecules, both are isoelectronic, meaning they contain a collective total of 14 electrons.
The nitrogen molecule, \(\mathrm{N}_{2}\), gets its electrons from two nitrogen atoms, each contributing 7 electrons. On the other hand, carbon monoxide, \(\mathrm{CO}\), has its electrons shared between carbon (6 electrons) and oxygen (8 electrons).
Despite having identical numbers of electrons, \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) exhibit different properties due to other factors, such as the types of bonds and the atoms involved. This brings an intriguing insight into how electron count doesn't solely dictate molecular properties, as these molecules show us differences in their stability and reactivity.

Electronegativity

Electronegativity is the measure of an atom's ability to attract and hold electrons within a bond. It plays a vital role in determining molecular structure and behavior, especially in the context of molecular orbital theory.
In molecules like \(\mathrm{CO}\), carbon and oxygen have different electronegativities, with oxygen being significantly more electronegative than carbon. This creates a polar bond within the molecule, making one end slightly negative (the oxygen end) and the other slightly positive (the carbon end).
Contrast this with \(\mathrm{N}_{2}\), where both nitrogen atoms have similar electronegativities, allowing for equal sharing of electrons within the molecule and resulting in a non-polar bond.

• This influence of electronegativity gives rise to different molecular properties.
• The \(\sigma_{2p}\) orbital in \(\mathrm{CO}\) is more stabilized compared to \(\mathrm{N}_{2}\) due to oxygen's higher electronegativity.

Understanding electronegativity helps in predicting the molecular orbital diagrams and resulting molecular behavior.

Triple bond

A triple bond is a type of chemical bond where three pairs of electrons are shared between two atoms. This kind of bond is significant in determining the strength and characteristics of a molecule.
Both \(\mathrm{N}_{2}\) and \(\mathrm{CO}\) have a triple bond, but the similarity largely ends there, as the differences in their atomic constituents greatly influence the characteristics of those bonds.

• In \(\mathrm{N}_{2}\), the triple bond is formed between the two nitrogen atoms.
• This bond is extremely strong and contributes to the overall stability and high bond dissociation energy of the molecule.
• As a result, \(\mathrm{N}_{2}\) is quite inert under normal conditions.

For \(\mathrm{CO}\), though the bond is formally classified as a triple bond, the difference in electronegativity between carbon and oxygen makes the bond polar.
This polarity can affect the molecule's reactivity and interaction with other substances, being fundamentally different from the triple bond in \(\mathrm{N}_{2}\).
Studying these bonds highlights that identical bond orders can manifest uniquely when atomic compositions change.

Molecular structure

Molecular structure is a term that describes the arrangement of atoms within a molecule, dictating the molecule's geometry and impacting its chemical and physical properties.
The molecular structure of a molecule is influenced by factors such as the presence of bonds, types of atoms, and their electronegativities. In the context of molecular orbital (MO) theory:

• \(\mathrm{N}_{2}\) has a symmetrical molecular structure with both nitrogen atoms contributing equally to the electron cloud, resulting in a non-polar diatomic molecule.
• Its stability and inertness are attributed to this symmetric structure and equal electronegativity between the nitrogen atoms.

On the other hand, \(\mathrm{CO}\)'s molecular structure is affected by the disparate electronegativities of carbon and oxygen atoms.
This contributes to an asymmetric electron distribution, leading to a polar molecule.
Such structural differences are critical, as they define how the molecules will interact with other substances, influencing their reactivity and applications in real-world contexts. Ultimately, understanding the interplay of isoelectronic configurations and electronegativity in shaping molecular structure enriches our grasp of molecular science.