Question

Choose the incorrect answer: (a) Bonding MO are formed by the additive effect of the atomic orbitals. (b) The electron density in the internuclear region in bonding MO is very low. (c) Energy of the antibonding MO is higher than that of the combining atomic orbitals. (d) Antibonding MO are represented by \(\sigma^{*}, \pi^{*}, \delta^{*}\), etc.

Short answer

Option (b) is incorrect. The electron density in the internuclear region in bonding Molecular Orbitals is high, not low.

Step-by-step solution

Understanding of Bonding MO

It is known that Bonding Molecular Orbitals (MO) are formed by the additive (or constructive) effect of atomic orbitals. So, option (a) is correct.

Understanding Electron Density in MO

Option (b) states that the electron density in the internuclear region in bonding MO is very low. This is incorrect. In bonding MO, the electron density in the internuclear region is usually high and not low. This high electron density causes attractive forces between the atomic nuclei.

Antibonding MO Energy Levels

Option (c) states that the energy of the antibonding MO is higher than that of the combining atomic orbitals. This is correct. Antibonding MO results from the subtractive effect (or destructive interference) of atomic orbitals, and they have higher energy than the original atomic orbitals.

Representation of Antibonding MO

Option (d) suggests that Antibonding MO are typically represented by \(\sigma^{*}, \pi^{*}, \delta^{*}\) etc. This is also correct. The presence of the star (*) denotes an antibonding orbital.

Key concepts

Bonding MO

Molecular orbitals (MO) are a fundamental concept in the study of chemical bonding. Bonding molecular orbitals, or bonding MOs, are of particular importance because they hold the electrons that allow atoms to bind together in molecules. Formed by the constructive interference of two atomic orbitals, bonding MOs have a lower energy than the individual atomic orbitals from which they originate. The result of this combination is an electron density that is increased between the two atoms, leading to a stable configuration.

When visualizing these orbitals, imagine two overlapping clouds of electron density that merge together enhancing the space between the nuclei. This region of increased electron density leads to an attractive force that holds the nuclei together and is a key feature of a chemical bond.

Antibonding MO

In contrast to bonding MOs, antibonding molecular orbitals (antibonding MOs) are characterized by higher energy and instability. They are formed when atomic orbitals combine in such a way that leads to the destructive interference of their wave functions. An easy way to spot an antibonding MO is by its notation, which includes an asterisk, like \( \sigma^{*} \) or \( \pi^{*} \).

These orbitals have a node, an area of zero electron density, between the two nuclei. The presence of a node decreases electron density where the two nuclei would experience the greatest attraction, thus weakening the bond. Filling an antibonding MO with electrons can negate the stabilizing effects of a bonding MO, potentially leading to the dissociation of a molecule if too many antibonding orbitals are occupied.

Electron Density in MO

The concept of electron density in molecular orbitals is instrumental in understanding chemical bonds. In a bonding MO, electron density is concentrated in the region between atomic nuclei, which is a favorable configuration leading to a bond formation. It's here where attractive forces overcome repulsion and enable nuclei to stay close together.

Conversely, in an antibonding MO, the electron density is lowest between the nuclei and places them on either side of a nodal plane. This distribution of electrons results in a repulsive interaction, which tends to push the nuclei apart, weakening or even preventing a bond. The understanding of electron density distribution is essential in predicting the stability and reactivity of molecules.

Orbital Hybridization

Another critical aspect of molecular bond theory is orbital hybridization. This occurs when atomic orbitals mix to form new, hybridized orbitals that influence the geometry and bonding properties of molecules. The classic types of hybrid orbitals are sp, sp2, and sp3, which describe the mixing of one s orbital with one, two, or three p orbitals, respectively.

The concept of hybridization helps explain why the molecular geometry of methane (CH4) is tetrahedral - its carbon atom has sp3 hybrid orbitals that arrange themselves as far apart as possible to minimize repulsion. Understanding the types and geometries of hybrid orbitals aids in predicting the shapes of molecules and their potential interactions with other substances.