Question

Use formal charge arguments to explain why \(\mathrm{CO}\) has a much smaller dipole moment than would be expected on the basis of electronegativity.

Short answer

The dipole moment of carbon monoxide (CO) is smaller than expected based on electronegativity due to the presence of formal charges in the molecule. The CO molecule has a triple bond with a formal charge of -1 on the carbon atom and +1 on the oxygen atom. These formal charges cause electron density redistribution, counteracting the electronegativity difference and resulting in a smaller dipole moment.

Step-by-step solution

Understanding dipole moments and electronegativity

A dipole moment is a measure of the separation of positive and negative charges in a molecule. The larger the difference in electronegativity between the two atoms in a bond, the greater the dipole moment. Electronegativity is the ability of an atom to attract electron density towards itself.

Analyzing formal charges in CO

To analyze formal charges, we need to consider the Lewis structure of the CO molecule. In the case of CO, it has a triple bond between the carbon and oxygen atoms and one lone pair of electrons on the carbon atom. The formal charges in a molecule are determined by subtracting the number of valence electrons an atom has initially from the number of electrons it has in its bonds and lone pairs in the final structure. Formal charge for Carbon (C): \(4 - 3(bonds) - 2(lone\ pair\ electrons) = -1\) Formal charge for Oxygen (O): \(6 - 3(bonds) - 4(lone\ pair\ electrons) = +1\)

Evaluating the effect of formal charges on dipole moment

The presence of formal charges in the CO molecule causes electron density to be shared more evenly between the carbon and oxygen atoms. This is because the negative formal charge on carbon and positive formal charge on oxygen cause an electron density redistribution that counteracts the difference in electronegativity. As a result, CO has a smaller dipole moment than would be expected based on the electronegativity differences between the carbon and oxygen atoms. In conclusion, using formal charge arguments, we can explain why the dipole moment of carbon monoxide (CO) is much smaller than what would be expected based on the electronegativity of carbon and oxygen atoms. The presence of formal charges in CO helps to redistribute electron density between the two atoms, reducing the expected dipole moment.

Key concepts

Dipole Moments

Dipole moments represent the distribution of electrical charges within a molecule. They occur due to the differences in electronegativity between two atoms forming a bond, where the more electronegative atom attracts more electrons towards itself. A molecule with a significant difference in electronegativity between its atoms will have a strong dipole moment. This is often visualized as an arrow, with the head pointing towards the negative charge and the tail at the positive charge.
When considering carbonate monoxide (CO), it may seem contradictory that its dipole moment isn't as pronounced as one would expect, given the electronegativity difference between carbon and oxygen. Despite oxygen being more electronegative, formal charge arguments help decipher this discrepancy. The unexpected dipole moment in CO reveals that there's more to the story than mere electronegativity—a fascinating illustration of how molecule structure impacts properties like polarity.

Electronegativity

Electronegativity is a concept in chemistry that describes an atom's tendency to attract a shared pair of electrons—or electron density—within a covalent bond. Atoms with high electronegativity, such as oxygen, fluorine, and nitrogen, pull electrons towards themselves more effectively than lesser electronegative atoms like carbon or hydrogen.
This electronegative difference between atoms in a molecule can lead to polar bonds, where one atom has a slight negative charge and the other a slight positive charge. Understanding electronegativity is crucial not only in predicting molecule behavior but also in explaining phenomena such as the reduced dipole moment in carbon monoxide, despite the clear difference in electronegativity between carbon and oxygen.

Lewis Structures

Lewis structures are a staple tool in chemistry for visualizing the bonds between atoms in molecules, as well as any lone pairs of electrons that may exist. They provide a diagrammatic representation that helps in determining the molecular geometry, the distribution of electrons, and the prediction of molecular properties.
The creation of a Lewis structure follows specific rules to allocate valence electrons correctly: counting the total valence electrons available, pairing electrons to form bonds, and filling non-bonding outer shells. For molecules like CO, Lewis structures are particularly informative. They detail how the triple bond forms and how the lone pairs contribute to the molecule's overall structure and charge distribution. A thorough understanding of Lewis structures is instrumental in comprehending more nuanced concepts such as formal charges and their influence on properties like dipole moments.

Formal Charges Calculation

The formal charge of an atom in a molecule is calculated by taking the number of valence electrons that the atom owns and subtracting the number of electrons it's assigned in the Lewis structure, both in bonds and as lone pairs.

1. Count the valence electrons the atom has in the unbound state.
2. Count the electrons surrounding the atom in the molecule, assigning one electron for each bond and counting lone electrons normally.
3. The formal charge is the difference between these counts.

To calculate this for CO, one must consider that carbon, typically having four valence electrons, and oxygen, with six, are assigned a differential number of electrons due to the triple bond and lone pairs. With a formal charge of -1 on carbon and +1 on oxygen in CO, the electron density gets distributed in a way that diminishes the dipole moment, balancing out the effect caused by the difference in electronegativity. The formal charge isn't the actual charge but a bookkeeping tool that, when combined with understanding of electronegativity and Lewis structures, allows chemists to understand the molecular behavior more thoroughly.