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: \(\mathrm{M}-\mathrm{C} \equiv \mathrm{O}\) 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. \(9.42\) and \(9.43 .\) 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

In summary, the carbon atom in carbon monoxide (CO) is more likely to form bonds to metals based on its lower electronegativity, the formal charges of the atoms, and the MO model which places more electron density near the carbon atom in the antibonding molecular orbitals.

Step-by-step solution

a. Electronegativity Analysis

Recall that electronegativity is the tendency of an atom to attract electrons towards itself. Based on this, we would expect the atom with lower electronegativity to form bonds to metals. Using the periodic table, we can find the electronegativities of carbon and oxygen: - Carbon (C) has an electronegativity of 2.55 - Oxygen (O) has an electronegativity of 3.44 Since carbon has a lower electronegativity than oxygen, we would expect the carbon atom to form bonds to metals.

b. Assigning Formal Charges

To assign formal charges to the atoms in CO, we will follow these steps: 1. Determine the number of valence electrons for each atom. 2. Determine the number of electrons assigned to each atom in the Lewis structure. 3. Calculate the formal charge for each atom using the formula: Formal Charge = Valence electrons - Assigned electrons For carbon: 1. Carbon has 4 valence electrons (from remaining in Group 4 of the periodic table). 2. In the Lewis structure of CO, carbon is triple-bonded to oxygen and has one lone pair. As there are 2 electrons in each bond and 2 electrons in the lone pair, a total of 6 electrons are assigned to carbon. 3. The formal charge of carbon is: 4 - 6 = -2 For oxygen: 1. Oxygen has 6 valence electrons (from remaining in Group 16 of the periodic table). 2. In the Lewis structure of CO, oxygen is triple-bonded to carbon, so it is assigned a total of 6 electrons (3 bonds * 2 electrons per bond = 6 electrons). 3. The formal charge of oxygen is: 6 - 6 = 0 Based on formal charges, carbon has a charge of -2, while oxygen has a charge of 0. As a more negatively charged atom, carbon would be more likely to form bonds to metals.

c. MO Model Prediction

In the MO model, the bonding molecular orbitals have more electron density near the more electronegative atom, while the antibonding molecular orbitals have more electron density near the less electronegative atom. Since we already determined that carbon is the less electronegative atom, antibonding MOs place more electron density near the carbon atom. This suggests that carbon should be the atom in CO that forms bonds to metals based on the MO model. In conclusion, based on electronegativity, formal charges, and the MO model, we predict that the carbon atom in carbon monoxide (CO) is the one that forms bonds to metals.

Key concepts

Understanding Electronegativity in ChemicalBonding

Electronegativity is a fundamental chemical concept that helps predict the nature and strength of bonds between atoms in molecules. It refers to the tendency of an atom to attract shared electrons in a chemical bond.

Elements across the periodic table have different electronegativities, with fluorine being the most electronegative and francium the least. Generally, the further to the right and closer to the top of the periodic table an element is, the higher its electronegativity.

Exploring the carbon monoxide molecule, with carbon at an electronegativity of 2.55 and oxygen at 3.44, it becomes apparent that oxygen has a stronger pull on the shared electrons. This is essential to understanding why, contrary to expectation, carbon is the atom that bonds to metals rather than oxygen. Despite carbon's lower electronegativity, which would usually make it less likely to bond with metals, the structure and distribution of electrons in CO prompt carbon to facilitate this bonding.

Understanding electronegativity also helps explain the polarity of molecules, reactivity, and many other chemical properties, making it an indispensable tool in predicting molecular behavior.

The Role of Formal Charges in Molecule Stability

Formal charges offer insight into the distribution of electrons within a molecule and help us predict the stability of molecular structures. They are calculated by balancing the number of valence electrons against the electrons assigned to an atom in a molecule.

To calculate the formal charge, you use the formula:

• Formal Charge = (Valence electrons) - (non-bonding electrons + 1/2 bonding electrons)

For instance, in carbon monoxide (CO), the calculations of formal charges show that the carbon atom carries a negative charge while the oxygen atom is neutral. Although this does not reflect the actual charge distribution due to resonance and molecular geometry, it aids in determining how a molecule might interact with other entities, such as metal ions.

Consequently, the carbon atom with a negative formal charge in CO is more likely to seek stabilization through bonding with positively charged metals—another piece of the puzzle explaining CO's bonding behavior.

Deciphering Molecule Interactions with the Molecular Orbital Model

The Molecular Orbital (MO) model is a sophisticated method that describes the electronic structure of molecules. It provides a way to visualize how atomic orbitals combine to form molecular orbitals, which can be either bonding or antibonding.

Bonding molecular orbitals result in a lower energy state and increased stability of a molecule, often localizing more electron density around the more electronegative atom. In contrast, antibonding molecular orbitals tend to concentrate electron density around the less electronegative atom and are associated with higher energy and less stable states.

In the context of the exercise, the MO model predicts that carbon would be the atom in CO to form bonds with metals since the carbon atom associates with the antibonding molecular orbitals. This conclusion aligns with both the electronegativity and formal charge assessments for carbon monoxide, providing a robust explanation for its bonding tendencies toward metals, despite the individual electronegativity of carbon itself.