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.

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#

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: \( \text{Formal charge}_{C} = 4 - \left(0+ \frac{6}{2}\right) = 4 - 3 = +1 \) 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): \( \text{Formal charge}_{O} = 6 - \left(4 + \frac{6}{2}\right) = 6 - 7 = -1 \) 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#

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.