Question

Consider three molecules: \(\mathrm{A}, \mathrm{B}\), and \(\mathrm{C}\). Molecule A has a hybridization of \(s p^{3}\). Molecule B has two more effective pairs (electron pairs around the central atom) than molecule A. Molecule C consists of two \(\sigma\) bonds and two \(\pi\) bonds. Give the molecular structure, hybridization, bond angles, and an example for each molecule.

Short answer

Molecule A: Tetrahedral structure, \(sp^3\) hybridization, bond angle of 109.5 degrees. Example: Methane (CH4). Molecule B: Octahedral structure, \(sp^3d^2\) hybridization, bond angle of 90 degrees. Example: Sulfur hexafluoride (SF6). Molecule C: Trigonal planar structure (one double bond and two single bonds), \(sp^2\) hybridization, bond angle of 120 degrees. Example: Ethene (C2H4).

Step-by-step solution

Molecule A: Hybridization \(sp^3\)

Molecular structure: Tetrahedral Hybridization: \(sp^3\) Bond angles: 109.5 degrees Example: Methane (CH4)

Molecule B: 2 more effective pairs than A

Since molecule A has a hybridization of \(sp^3\) with 4 effective pairs (electron pairs around the central atom), molecule B will have 6 effective pairs. This implies that molecule B has an octahedral structure. Molecular structure: Octahedral Hybridization: \(sp^3d^2\) Bond angles: 90 degrees Example: Sulfur hexafluoride (SF6)

Molecule C: 2 \(\sigma\) bonds and 2 \(\pi\) bonds

Having two \(\sigma\) bonds and two \(\pi\) bonds indicates that the molecule is a double bond. A compound with one double bond typically has a hybridization of \(sp^2\). Molecular structure: Trigonal planar (one double bond and two single bonds) Hybridization: \(sp^2\) Bond angles: 120 degrees Example: Ethene (C2H4)

Key concepts

Hybridization

Hybridization is an essential concept in chemistry that explains how atomic orbitals mix to form new hybrid orbitals. These hybrid orbitals can explain the shape and geometry of a molecule.
In the case of Molecule A from the exercise, it has an \( sp^3 \) hybridization. This means it combines one s orbital and three p orbitals to form four equivalent \( sp^3 \) hybrid orbitals. These orbitals arrange themselves in a tetrahedral geometry, which minimizes repulsion between the electron pairs.
For Molecule B, which has a hybridization of \( sp^3d^2 \), there are six orbitals involved. Here, one s orbital, three p orbitals, and two d orbitals mix, leading to an octahedral structure. Such structures are common in molecules that have more electron pairs surrounding their central atom.
Molecule C, with its \( sp^2 \) hybridization, combines one s orbital and two p orbitals. This results in three \( sp^2 \) hybrid orbitals with a trigonal planar arrangement. This is common in molecules with double bonds, as they involve both \( \sigma \) and \( \pi \) bonds.

Bond Angles

Bond angles are the angles formed between adjacent bonds of a molecule. Understanding them helps to visualize the molecule's geometry.
In Molecule A, with a tetrahedral structure, the bond angles are exactly 109.5 degrees. This specific angle is due to the four \( sp^3 \) hybrid orbitals being placed at the corners of a tetrahedron, providing a uniform distribution.
With Molecule B's octahedral configuration, the bond angles between adjacent pairs of electron pairs are 90 degrees. This is because the \( sp^3d^2 \) hybrid orbitals position themselves in a symmetrical octahedron, optimizing distance and minimizing repulsion.
For Molecule C, which is trigonal planar, the bond angles are approximately 120 degrees. This geometry arises from the \( sp^2 \) hybridization, ensuring the planar arrangement of electron pairs and involvement of \( \sigma \) and \( \pi \) bonds.

Electron Pairs

Electron pairs significantly influence a molecule's shape and hybridization. They repel each other, and this repulsion dictates the geometry.
In Molecule A, with a tetrahedral shape, there are four electron pairs around the central atom. These could be bonded pairs or lone pairs, creating the \( sp^3 \) hybrid formation.
Molecule B, on the other hand, features more complexity with six effective electron pairs. These pairs lead to an octahedral structure with \( sp^3d^2 \) hybridization, which is typical for molecules like sulfur hexafluoride (SF₆). The arrangement minimizes repulsion by orienting pairs at 90 degrees.
Molecule C contains distinctions as well. Here, the focus is on both \( \sigma \) bonds and \( \pi \) bonds, reflected in its three electron pair layout. Although less than Molecule A and B, the electron pair arrangement allows for \( sp^2 \) hybridization, ideal for forming double bonds.

Sigma and Pi Bonds

\( \sigma \) and \( \pi \) bonds are crucial for understanding bond types and strengths in molecules. These bonds arise from different types of atomic orbital overlaps.
\( \sigma \) bonds are single bonds formed by the direct overlap of orbitals. They're strong and possess head-on overlap, providing primary structural stability in molecules. Molecule C has two \( \sigma \) bonds, accounting for its formation of double bonds along with \( \pi \) bonds.
\( \pi \) bonds, by contrast, are formed when parallel p orbitals overlap side-by-side. They only occur in addition to \( \sigma \) bonds, meaning without an initial \( \sigma \) bond, a \( \pi \) bond cannot exist. In Molecule C, the two \( \pi \) bonds create the additional bond layers found in double bonds like those in ethene (C₂H₄). These bonds affect the rigidity and planarity of a molecule, often dictating reactivity and interaction potentials.