Question

Carbon monoxide (CO) forms bonds to a variety of metals and metal ions. Its ability to bond to iron in hemoglobin is the reason that $\mathrm{CO}$ is so toxic. The bond carbon monoxide forms to metals is through the carbon atom: $$\mathbf{M}-\mathbf{C}\equiv \mathbf{O}$$ a. On the basis of electronegativities, would you expect the carbon atom or the oxygen atom to form bonds to metals? b. Assign formal charges to the atoms in CO. Which atom would you expect to bond to a metal on this basis? c. In the MO model, bonding MOs place more electron density near the more electronegative atom. (See the HF molecule in Figs. $4-54$ and $4-55.$ ) Antibonding MOs place more electron density near the less electronegative atom in the diatomic molecule. Use the MO model to predict which atom of carbon monoxide should form bonds to metals.

Short answer

Based on electronegativity, the less electronegative carbon atom is expected to form bonds with metals. Using formal charges, the positively charged carbon atom is more likely to bond to a metal. In the MO model, the carbon atom of carbon monoxide has a higher electron density in the antibonding orbitals, making it more likely to bond with metals. Overall, the carbon atom is more likely to form bonds with metals in CO.

Step-by-step solution

Compare Electronegativities of Carbon and Oxygen

To determine which atom is more likely to form bonds with metals based on electronegativity, we must look at the electronegativities of carbon and oxygen. Oxygen has an electronegativity of 3.44, while carbon has an electronegativity of 2.55. The higher the electronegativity, the greater the attraction an atom has for electrons. Oxygen is more electronegative than carbon, meaning it attracts electrons more strongly. Thus, we might expect the less electronegative carbon atom to form bonds with metals, as metal atoms generally have lower electronegativity and tend to lose electrons. #b. Formal Charges Calculation#

Calculate Formal Charges

Formal charges of atoms in a molecule can help predict which atom is more likely to form bonds. The individual formal charges can be calculated based on the number of valence electrons, the number of electrons in lone pairs, and the number of electrons in bonds assigned to a specific atom. For carbon monoxide (CO), we have: Carbon formal charge: $\text{Formal charge}=\text{valence electrons}-(\text{non-bonding electrons}+\frac{\text{bonding electrons}}{2})$ Carbon has 4 valence electrons. There are no non-bonding electrons, and there are 6 bonding electrons (triple bond). So, the formal charge of carbon is: ${\text{Formal charge}}_C=4-(0+\frac{6}{2})=4-3=+1$ For oxygen: Oxygen has 6 valence electrons, and in the case of CO, it has 2 non-bonding pairs (4 non-bonding electrons). The number of bonding electrons is the same as for carbon (6 electrons): ${\text{Formal charge}}_O=6-(4+\frac{6}{2})=6-7=-1$ Taking formal charges into account, we would expect the positively charged carbon atom to bond to a metal, which often carries a positive charge. #c. MO Model Prediction#

Use the MO Model to Predict Bond Formation

In the molecular orbital (MO) model, bonding orbitals place more electron density around the more electronegative atom (oxygen) in diatomic molecules, while the antibonding orbitals place more electron density around the less electronegative atom (carbon). Since metals form bonds when their atoms' electron density overlaps with that of another species, which in this case, is CO, it is likely that the less electronegative carbon atom with more electron density in antibonding orbitals will readily interact and bond with metal atoms. Thus, using the MO model, we predict that the carbon atom of carbon monoxide should form bonds to metals.

Key concepts

Electronegativity

Electronegativity is a measure of how strongly an atom can attract electrons towards itself. In the context of carbon monoxide (CO), we compare the electronegativities of carbon and oxygen to understand their bonding behavior.
Oxygen, having an electronegativity of 3.44, is more electronegative than carbon, which has an electronegativity of 2.55. This means that oxygen has a stronger pull on electrons.
However, when CO forms bonds with metals, it is the carbon atom that typically forms the bond, despite the difference in electronegativity.

• Metals generally have a lower electronegativity.
• The less electronegative carbon can more easily share its electrons with metals.

This is why, even though oxygen attracts electrons more strongly, carbon is the atom that bonds with metals.

Formal Charge

Formal charge helps decide which atom in a molecule is more likely to bond with another species. It’s calculated by considering the number of valence electrons, lone pairs, and bonds on an atom.
For carbon monoxide, formal charges give insight into its bonding characteristics:

• Carbon has a formal charge of +1.
• Oxygen has a formal charge of -1.

Formal charge simplifies to:$$\text{Formal charge}=\text{Valence electrons}-(\text{Non-bonding electrons}+{\displaystyle \frac{\text{Bonding electrons}}{2}})$$With carbon's formal charge being +1, it tends to bond with positively charged metals, making it the preferred bonding site. Despite oxygen being more electronegative, the formal charge distribution favors carbon in metal interactions.

Molecular Orbital Theory

Molecular Orbital (MO) Theory provides a deeper insight into chemical bonding by describing the behavior of electrons in a molecule. In MO theory, electrons in a molecule are not restricted to individual atoms but are distributed over the entire molecule's orbitals.
In CO, MO Theory suggests the following:

• Bonding MOs tend to have more electron density around the more electronegative atom, oxygen.
• Antibonding MOs have more electron density around the less electronegative atom, carbon.

Although oxygen attracts electrons more effectively, the overlap between metal orbitals and CO's antibonding orbitals makes carbon a preferable site for bonding. This results from its electron density characteristics that facilitate interaction and bonding with metal atoms.

Metal Bonding

Metal bonding with carbon monoxide is a fascinating process. CO can bond with metals like iron, playing an essential role in biological systems such as hemoglobin.
When CO bonds with metals, the electron-sharing involves the carbon atom. This interaction is crucial as carbon, despite being less electronegative, possesses the necessary electron cloud that facilitates bonding.
Key points regarding CO-metal bonding include:

• The metal bonds through a coordinate covalent bond, where CO donates a pair of electrons.
• The bonding often involves the d-orbitals of the metal, which readily accepts electrons from CO.
• Such bonds influence industrial catalysis and biological metal-ligand interactions.

The carbon atom's ability to engage in metal bonding significantly impacts the reactivity and functionality of metallic systems.