Use formal charge arguments to explain why CO has a much smaller dipole moment than would be expected on the basis of electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The dipole moment is a measure of the polarity of a molecule, which depends on the difference in electronegativity between the two bonded atoms and the distance between them. Formal charge is the charge assigned to each atom in a molecule assuming that the bonding electrons are evenly shared between the two atoms.

To determine the expected dipole moment in CO based on electronegativity values, we need to consider the electronegativity difference between the carbon (C) and oxygen (O) atoms. The electronegativity values for carbon and oxygen are 2.55 and 3.44 (Pauling scale), respectively. The difference in electronegativity is: Electronegativity difference = \(\lvert3.44 - 2.55 \rvert\) = 0.89 By comparing this difference to other molecules, we could estimate that the dipole moment of CO should be significant due to the considerable electronegativity difference between carbon and oxygen atoms.

Now, we have to calculate the formal charge for each atom in CO. The formal charge for an atom can be calculated using the following formula: Formal charge = (Valence electrons) - (Non-bonding electrons) - (Bonding electrons / 2) For carbon (C): Valence electrons: 4 Non-bonding electrons: 0 Bonding electrons: 10 (as carbon - triple bond - oxygen) Formal charge on C = \(4 - 0 - (10/2) = -1\) For oxygen (O): Valence electrons: 6 Non-bonding electrons: 4 Bonding electrons: 10 (as carbon - triple bond - oxygen) Formal charge on O = \(6 - 4 - (10/2) = +1\)

The formal charges on carbon (-1) and oxygen (+1) indicate that there is a significant electron density shift toward the oxygen atom, which could lead to the expectation of a significant dipole moment. However, the actual dipole moment of CO is smaller than what is expected. The formal charge argument helps to explain this observation by considering the bonding orbitals between the carbon and oxygen atoms in CO. The triple bond between C and O consists of one σ-bond and two π-bonds, formed by overlapping hybridized orbitals from both atoms. The two π-bonds in the triple bond might have a significant delocalization of electron density, reducing the polarity of the molecule. This delocalization of electron density dampens the expected polarity effect due to the electronegativity difference between carbon and oxygen. Consequently, CO has a smaller dipole moment than expected based solely on electronegativity.