Question

For each of these contour representations of molecular orbitals, identify (a) the atomic orbitals (s or \(p\) ) used to construct the MO (b) the type of MO ( \(\sigma\) or \(\pi\) ), (c) whether the MO is bonding or antibonding, and (d) the locations of nodal planes. [Sections \(9.7\) and 9.8]

Short answer

Example 1: σ bonding between p orbitals: (a) Atomic orbitals: Two \(p\) orbitals from adjacent atoms. (b) Type of MO: \(\sigma\) molecular orbital. (c) Bonding or antibonding: Bonding. (d) Nodal planes: One nodal plane, perpendicular to the internuclear axis and through the nuclei. Example 2: π bonding between p orbitals: (a) Atomic orbitals: Two \(p\) orbitals from adjacent atoms. (b) Type of MO: \(\pi\) molecular orbital. (c) Bonding or antibonding: Bonding. (d) Nodal planes: Two nodal planes - one along the internuclear axis and the other perpendicular to the axis and through the nuclei.

Step-by-step solution

Example 1: σ bonding between p orbitals

(a) Atomic orbitals: In this example, the atomic orbitals involved are two \(p\) orbitals from adjacent atoms. (b) Type of MO: The overlap between the lobes of the two p orbitals along the internuclear axis results in a \(\sigma\) molecular orbital. (c) Bonding or antibonding: Since the lobes of the p orbitals on each atom are in-phase, the resulting molecular orbital is bonding. (d) Nodal planes: For a \(\sigma\) molecular orbital formed by the overlap of p orbitals, there is only one nodal plane which is perpendicular to the internuclear axis and goes through the nuclei.

Example 2: π bonding between p orbitals

(a) Atomic orbitals: In this example, the atomic orbitals involved are also two \(p\) orbitals from adjacent atoms. (b) Type of MO: The overlap between lobes of the two p orbitals parallel to the internuclear axis results in a \(\pi\) molecular orbital. (c) Bonding or antibonding: Since the lobes of the p orbitals on each atom are in-phase, the resulting molecular orbital is bonding. (d) Nodal planes: For a \(\pi\) molecular orbital formed by the overlap of p orbitals, there are two nodal planes. One is along the internuclear axis, separating the lobes of each p orbital, and the other is perpendicular to the internuclear axis and goes through the nuclei.

Key concepts

Sigma Bonding

In chemistry, sigma (\(\sigma\)) bonding is a type of covalent bond where the electron density is concentrated directly along the bond axis between two atomic nuclei. This occurs through the head-on overlap of atomic orbitals.

• A \(\sigma\) bond can form between any combination of \(s\) or \(p\) orbitals, as long as their overlap along the internuclear axis is maximum.
• In the case of \(s\) orbitals, they overlap directly because of their spherical shape.
• For \(p\) orbitals, the lobes point toward each other and overlap head-on.

The result of a \(\sigma\) bond is not just stronger bonding, but also a more stabilized molecular orbital due to the direct overlap maximizing shared electron density. Since this overlap is symmetrical about the bond axis, \(\sigma\) bonds are typically single bonds, serving as the primary bond between atoms.

Pi Bonding

Pi (\(\pi\)) bonding is a distinct form of molecular bonding that complements \(\sigma\) bonds. Unlike \(\sigma\) bonds that occur along the internuclear axis, \(\pi\) bonds arise from the parallel overlap of \(p\) orbitals.

• In a \(\pi\) bonding, electron density is localized above and below the axis connecting the two nuclei.
• Pi bonds are often found in double or triple bonds, where they add additional bonding strength to the \(\sigma\) bond already present.
• Due to their nature of overlap, \(\pi\) bonds prevent the free rotation of the bonded atoms, thereby affecting molecule geometry.

While \(\pi\) bonds are generally weaker than \(\sigma\) bonds due to the less effective side-to-side overlap, they are crucial for the rigidity and structure of many compounds, such as those containing double bonds.

Nodal Planes

Nodal planes are theoretical planes or spaces within molecular orbitals where the probability of finding an electron is zero. These are essential for understanding molecular shapes and bonding types.

• In molecular orbitals, nodal planes arise due to the wave nature of electrons, which can create nodes as a result of destructive interference.
• Sigma (\(\sigma\)) bonds typically have fewer nodal planes compared to \(\pi\) bonds because of their head-on overlap.
• For instance, a \(\sigma\) molecular orbital formed from \(p\) orbitals generally has one nodal plane, which is perpendicular to the bond axis.

In contrast, \(\pi\) bonds commonly exhibit at least two nodal planes, which influences the electron cloud distribution and molecular stability. Identifying nodal planes is important for predicting and visualizing molecular properties.

Atomic Orbitals

Atomic orbitals are regions within an atom where there's a high probability of finding electrons. Each type of atomic orbital has a distinctive shape and energy level which determines how these orbitals bond with others.

• The main types of atomic orbitals are \(s\), \(p\), \(d\), and \(f\), with each having a unique geometry.
• \(s\) orbitals are spherical, allowing uniform overlap for bonding.
• \(p\) orbitals are dumbbell-shaped, allowing for both head-on (for \(\sigma\) bonds) and side-by-side (for \(\pi\) bonds) overlap.

Understanding atomic orbitals is vital to comprehend how different atoms interact to form molecules. The shape and energy of these orbitals define how effectively they can overlap to form molecular orbitals.

Antibonding

Antibonding orbitals are formed when atomic orbitals combine in such a way that their electron wave functions destructively interfere. This leads to decreased electron density between the nuclei, effectively working against bond formation.

• Antibonding molecular orbitals are denoted with an asterisk, such as \(\sigma^*\) or \(\pi^*\).
• They have higher energy than bonding orbitals because the electron density is concentrated outside the region between the two nuclei.
• The presence of electrons in antibonding orbitals can weaken or prevent bond formation, sometimes leading to molecular instability.

In essence, these orbitals are critical in balancing out the electron distribution in a molecule, influencing its stability and reactionism. Recognizing the role of antibonding orbitals helps in understanding why certain molecules form while others do not.