Question

Carbon monoxide (CO) is a poisonous compound due to its ability to bind strongly to \(\mathrm{Fe}^{2+}\) in the hemoglobin molecule. The molecular orbitals of \(\mathrm{CO}\) have the same energy order as those of the \(\mathrm{N}_{2}\) molecule. (a) Draw a Lewis structure of \(\mathrm{CO}\) and assign formal charges. Explain why CO has a rather small dipole moment of \(0.12 \mathrm{D}\). (b) Compare the bond order of \(\mathrm{CO}\) with that from molecular orbital theory. (c) Which of the atoms \((\mathrm{C}\) or \(\mathrm{O})\) is more likely to form bonds with the \(\mathrm{Fe}^{2+}\) ion in hemoglobin?

Short answer

CO has a Lewis structure bond order of 3; C likely bonds with Fe²⁺ in hemoglobin.

Step-by-step solution

Drawing the Lewis Structure of CO

To draw the Lewis structure of carbon monoxide, we first count the total number of valence electrons. Carbon (C) has 4 valence electrons, and oxygen (O) has 6 valence electrons, giving a total of 10 valence electrons. We create a triple bond between carbon and oxygen, where each contributes three electrons to share. This uses up 6 electrons, leaving 2 lone pairs, one on carbon and one on oxygen. Assigning formal charges: Carbon will have a formal charge of -1, and oxygen will have a formal charge of +1.

Explaining Dipole Moment

The dipole moment of CO is rather small (0.12 D) because the molecule, although polar with formal charges, has a significant amount of electron sharing that reduces the electrostatic difference across the molecule. The small separation of formal charges and nearly equal electronegativity of C and O lead to a small dipole moment.

Calculating Bond Order using Lewis Structure

In the Lewis structure, CO has a triple bond, which gives a bond order of 3, as bond order is defined as the number of bonded electron pairs between two atoms.

Calculating Bond Order using Molecular Orbital Theory

According to molecular orbital theory and the energy order of CO being similar to N₂, CO has 10 valence electrons distributed in molecular orbitals. The filling order leads to fully filled bonding orbitals and partially filled antibonding orbitals. The bond order is calculated as 1/2(Bonding electrons - Antibonding electrons), resulting in a bond order of 3 since there are no antibonding electrons.

Determining Bond Formation with Fe^2+

Carbon is more likely to form bonds with the Fe²⁺ ion in hemoglobin. This is due to carbon having a lone pair of electrons which is less tightly held compared to oxygen, allowing carbon to donate it more easily to form a coordinate bond with Fe²⁺.

Key concepts

Lewis Structure

The Lewis structure of carbon monoxide (CO) is a useful way to visualize the arrangement of atoms and electrons in the molecule. First, we need to account for all valence electrons. Carbon, with 4 valence electrons, and oxygen, with 6, combine to yield a total of 10 electrons. In CO, a triple bond is formed between carbon and oxygen, using up 6 electrons. This means both carbon and oxygen will share three pairs of electrons, which gives it high bond strength.
The remaining 4 electrons create lone pairs: one on carbon and another on oxygen. This distribution leaves carbon with a formal charge of -1 and oxygen with a formal charge of +1. Even though oxygen might seem "happier" with more electrons due to its higher electronegativity, this arrangement allows for maximum bonding capacity.

Molecular Orbital Theory

Molecular orbital theory provides a different perspective on bonding compared to the Lewis structure representation. It considers how atomic orbitals combine to form molecular orbitals that can be bonding or antibonding. For CO, the molecular orbitals follow a similar order as those in nitrogen gas (N₂).
In CO, there are 10 valence electrons to fill these molecular orbitals. All electrons fill the bonding molecular orbitals without any in the antibonding orbitals, due to the symmetry and similar energy levels of carbon and oxygen. This results in a stable molecule with a bond order of 3. Without antibonding electrons, the molecule achieves an optimal bond strength.

• Bonding Molecular Orbitals: Electrons contribute positively to bond strength.
• Antibonding Molecular Orbitals: Usually weaken bonds if occupied, but none are filled in CO for this case.

Dipole Moment

The dipole moment of carbon monoxide is a measure of the overall polarity of the molecule, which is quite low at 0.12 Debye. Even with formal charges on carbon and oxygen, the near-equal electronegativity results in minimal net charge separation. This is because the shared electrons in the triple bond and lone pairs create a scenario where there is significant electron sharing.
The outcome is a weak dipole, despite the substantial difference in electron affinity typically found between carbon and oxygen. The dipole moment represents the lack of significant charge separation across the molecular span, manifesting in CO being almost nonpolar.

Bond Order

In chemistry, bond order is an important concept that indicates the number of chemical bonds between a pair of atoms. For CO, the bond order can be determined both by Lewis structures and molecular orbital theory. Both methods agree that the bond order is 3, which indicates a strong triple bond.
The triple bond consists of one sigma bond and two pi bonds. A higher bond order reflects greater bond strength and shorter bond length, correlating with the physical properties such as the stability of CO. Bond order not only gives insight into bond strength but also into potential reactivity and stability of the molecule.

Coordinate Bonding

Coordinate bonding, also known as dative bonding, occurs when one atom provides both electrons needed for a bond. In CO, this is significant when it interacts with metals like the Fe²⁺ ion in hemoglobin. Carbon is more likely to engage in coordinate bonding due to its less tightly held lone pair.
This lone pair on carbon can be donated to a metal like \(\text{Fe}^{2+}\), forming a robust coordinate complex. In biological systems, this makes CO a potent inhibitor because it binds stronger to iron in hemoglobin than oxygen, disrupting normal oxygen transport. Understanding this aspect of CO's chemistry unveils why it is so toxic and how it can outcompete oxygen in critical biological functions.