Question

Dipole moment of \(\mathrm{CO}\) is very small \((0.1 \mathrm{D})\) (inspite of the larger difference in electronegativity between \(\mathrm{C}\) and \(\mathrm{O}\) ) and negative end of the dipole lies near the less electronegative C-atom. It is due to: (a) Occupation of antibonding molecular orbitals (b) Occupation of the lone pairs as well as empty do bitals (c) Occupation of lone-pairs (d) Occupation of empty d-orbitals

Short answer

The correct answer is (c) Occupation of lone-pairs.

Step-by-step solution

Understanding the Dipole Moment

The dipole moment is a vector quantity indicating the separation of charge in a molecule. A greater difference in electronegativity typically leads to a higher dipole moment.

Examining the Molecule CO

Carbon monoxide (CO) has a small dipole moment of 0.1 D. Normally, oxygen, being more electronegative than carbon, should attract electrons more, resulting in a dipole moment towards O.

Identifying Charge Localization

The negative end of the dipole being near the carbon atom suggests a charge localization opposite of what is expected based on electronegativity alone.

Analyzing Orbital Occupations

This unusual dipole placement can be explained by the electronic configuration; the lone pair of electrons on carbon contributes to the charge distribution, localizing electron density more on carbon.

Key concepts

Carbon Monoxide (CO)

Carbon Monoxide, often abbreviated as CO, is a simple diatomic molecule consisting of one carbon atom and one oxygen atom. Despite its simplicity, the molecule is quite interesting due to its chemical properties. Oxygen is more electronegative than carbon, meaning it tends to attract electrons more strongly. However, what makes CO particularly peculiar is its dipole moment, which is very small at approximately 0.1 Debye.

Usually, we would expect the electron density to be pulled toward the oxygen atom due to its higher electronegativity. But in CO, the small dipole moment indicates a charge localization anomaly. In this molecule, the negative end of the dipole is closer to the carbon atom, not the oxygen. This unusual localization is due to how the molecule's electrons are distributed, influenced by the molecular orbitals they occupy.

Electronegativity

Electronegativity describes the ability of an atom in a molecule to attract shared electrons. The higher an atom's electronegativity, the stronger its pull on electrons. In the periodic table, electronegativity increases from left to right and decreases down a group. In carbon monoxide, oxygen's electronegativity is greater than carbon's, which you'd expect would lead to a dipole moment with the negative pole directed toward oxygen.

However, actual measurements reveal a very small dipole toward carbon instead. This counterintuitive observation can be explained by the molecular orbital theory. Electrons follow orbitals that are influenced by their energy states and how atoms hybridize in a molecule. Hence, electronegativity, while a valuable indicator, doesn't account for every aspect of electronic distribution.

Molecular Orbitals

Molecular orbital theory offers a deeper understanding of how electrons are distributed in molecules like carbon monoxide. In this theory, atomic orbitals combine to form molecular orbitals—these can be bonding, antibonding, or nonbonding, influencing electron distribution and, consequently, molecular properties.

In CO, the molecular orbitals result from the combination of carbon and oxygen's atomic orbitals. Some electrons occupy pi and sigma bonding orbitals, but crucially, the lone pair on carbon occupies a non-bonding molecular orbital. This lone pair plays a significant role in charge localization, causing the electron density to tilt more toward carbon, despite the oxygen's greater electronegativity.

• Bonding orbitals stabilize the molecule by holding atoms together.
• Antibonding orbitals work against bond formation.
• Non-bonding orbitals have no direct effect on bond strength but can affect properties like dipole moments.

Understanding molecular orbitals thus helps clarify why carbon monoxide exhibits such distinct properties.