Question

\(\pi\) bonds are formed by which of the following orbitals? a. Two s-orbitals b. Two \(p\) -orbitals c. One s- and one \(p\) -orbital d. Two sp \(^{2}\) -hybridized orbitals

Short answer

b. Two p-orbitals.

Step-by-step solution

Review Types of Orbitals

Understand what s, p, and sp^2 orbitals are. s-orbitals are spherical, p-orbitals are dumbbell-shaped, and sp^2 orbitals are a mix of s and p characteristics, forming a trigonal planar shape.

Identify \( \pi \)-Bond Formation

A \( \pi \)-bond is formed by the parallel overlap of two p-orbitals. This occurs when electrons occupy the same region of space above and below the nuclei of bonding atoms.

Evaluate Each Option

Evaluate each option: a. Two s-orbitals: Incorrect because s-orbitals can only form \sigma-bonds. b. Two p-orbitals: Correct because \( \pi \)-bonds form from p-orbitals overlapping. c. One s- and one p-orbital: Incorrect, this combination forms a \sigma-bond. d. Two sp^2-hybridized orbitals: Incorrect, these orbitals also form \sigma-bonds.

Select the Correct Answer

Based on the evaluation, the correct answer is option b.

Key concepts

Molecular Orbitals

When atoms come together to form molecules, their atomic orbitals combine to create molecular orbitals. These molecular orbitals are the regions in a molecule where electrons are likely to be found. Here are a few key points:

• Atomic orbitals (like s, p, and d) blend to form new orbitals as molecules form.
• Molecular orbitals can be bonding or antibonding.
• Bonding orbitals are lower in energy and help hold the atoms together.
• Antibonding orbitals are higher in energy and can weaken or prevent bonding.

For example, when two p-orbitals overlap side-by-side, they create a \( \text{pi} \)-bonding orbital and a \( \text{pi}^* \)-antibonding orbital. Each molecular orbital can hold a maximum of two electrons.

Pi-Bonds

Pi-bonds (\pi\text{-bonds}) are a type of covalent bond formed by the sideways overlap of p-orbitals. They are different from sigma (σ) bonds, which result from the head-on overlap of orbitals. Here are the essentials to know about \( \text{pi} \)-bonds:

• They occur in double and triple bonds; a double bond has one σ-bond and one \( \text{pi} \)-bond, while a triple bond has one σ-bond and two \( \text{pi} \)-bonds.
• \pi\text{-bonds} restrict the rotation of bonded atoms, giving molecules a fixed structure.
• These bonds contribute to the molecule's overall electron density above and below the plane of the atoms.

A classic example of \( \text{pi} \)-bonding happens in ethene (C2H4), where each carbon atom forms three σ-bonds and one \( \text{pi} \)-bond. This \( \text{pi} \)-bond results from the side-by-side overlap of unhybridized p-orbitals on each carbon.

P-Orbitals

P-orbitals have a distinctive dumbbell shape and are oriented at right angles to each other. Understanding p-orbitals is crucial for grasping how \( \text{pi} \)-bonds form. Here are some key points about p-orbitals:

• There are three p-orbitals (px, py, and pz) in each energy level above the first.
• Each p-orbital can hold up to two electrons with opposite spins.
• The shape and orientation allow for sideways (parallel) overlaps, essential for \( \text{pi} \)-bond formation.
• When two p-orbitals from different atoms overlap, it creates a \( \text{pi} \)-bond, contributing to the multiple bonds in molecules like alkenes and alkynes.

For instance, in a carbon-carbon double bond, one p-orbital from each carbon atom overlaps to form the \( \text{pi} \)-bond, while the other bonds are \( \text{sigma} \)-bonds formed by hybrid orbitals.