Question

Consider three molecules: A, \(\mathrm{B},\) and \(\mathrm{C}\) . Molecule \(\mathrm{A}\) has a hybridization of \(s p^{3} .\) Molecule \(\mathrm{B}\) 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 has an sp³ hybridization, a tetrahedral structure, a bond angle of 109.5°, and an example is Methane (CH₄). Molecule B has an sp³d² hybridization, an octahedral structure, a bond angle of 90°, and an example is Sulfur hexafluoride (SF₆). Molecule C has an sp² hybridization, a trigonal planar structure, a bond angle of 120°, and an example is Sulfur dioxide (SO₂).

Step-by-step solution

Molecule A

Since molecule A has an sp³ hybridization, it will have four effective pairs (electron pairs around the central atom). The molecular structure will be tetrahedral. The bond angle for an sp³ hybridized molecule is generally 109.5°. Example: Methane (CH₄), where the central atom (Carbon) has 4 sigma bonds with hydrogen atoms, resulting in an sp³ hybridization.

Molecule B

As molecule B has two more effective pairs than molecule A, it will have six effective pairs. Based on the electronic geometry, the molecular structure will be octahedral. For six effective pairs, the hybridization will be an sp³d² (s, three p, and two d orbitals). The bond angles in an octahedral molecule are typically 90°. Example: Sulfur hexafluoride (SF₆), where the central atom (Sulfur) has six sigma bonds with six fluorine atoms, resulting in an sp³d² hybridization.

Molecule C

As molecule C has two σ bonds and two π bonds, this points towards a double bond scenario. For such molecule, we have three effective pairs: one σ bond, one π bond, and one lone pair. The molecular structure will be trigonal planar. In this case, hybridization will be sp² (s, and two p orbitals). The bond angles in a trigonal planar molecule are typically 120°. Example: Sulfur dioxide (SO₂), where the central atom (Sulfur) has one sigma bond and one pi bond with each oxygen atom, resulting in an sp² hybridization.

Key concepts

Molecular Structure

When looking at the molecular structure of a molecule, it essentially describes the three-dimensional arrangement of atoms around a central atom. Depending on the number of effective pairs, or electron pairs, surrounding this central atom, different shapes or geometries emerge.
- **Tetrahedral Geometry:** This is the molecular structure you get with four effective pairs, as seen in molecule A. This arrangement places all four regions of electron density as far apart as possible, leading to a tetrapod shape. Methane (CH₄) is a classic example fitting this description. - **Octahedral Geometry:** For molecule B, having six effective pairs results in an octahedral structure. The symmetry allows all bonds to experience equal repulsion, maintaining the structure’s integrity. Sulfur hexafluoride (SF₆) exemplifies this shape. - **Trigonal Planar Geometry:** For molecule C with three effective pairs, the molecular structure becomes trigonal planar. This shape makes a flat, triangular arrangement, like in sulfur dioxide (SO₂).
Understanding these structures is crucial for predicting the properties and behaviors of different molecules, as the geometry influences both physical and chemical characteristics.

Bond Angles

Bond angles are the angles between adjacent bonds of atoms surrounding a central atom. These angles can be predicted based on the molecular structure and are influenced by the repulsions between electron pairs. - **Tetrahedral Bond Angle:** With a tetrahedral molecular structure in molecule A, the bond angles are approximately 109.5°. This is because the four bonds try to remain as far apart as possible in 3D space, minimizing repulsion. - **Octahedral Bond Angle:** For the octahedral geometry characteristic of molecule B, bond angles are a neat 90°. This shape neatly divides space around the central atom, maximizing distance between the six pairs. - **Trigonal Planar Bond Angle:** In molecule C, with its trigonal planar structure, the bond angles are around 120°. This flat structure spreads the bonds equally in a plane, balancing the repulsion and attraction forces.
Bond angles are fundamental to determining the molecule's shape and, by extension, its behavior and reactivity.

Effective Pairs

Effective pairs in chemistry refer to the total number of electron pair regions around a central atom. These include both bonding pairs (used in bonds) and lone pairs (non-bonding electrons). Their arrangement dictates the molecular shape and geometry. - **For Molecule A:** It has four effective pairs, aligning with an sp³ hybridization, indicative of a tetrahedral geometry. Its known structure is simple and stable. - **For Molecule B:** With six effective pairs, because it possesses two more than molecule A, we observe an octahedral geometry that’s symmetrical and optimal for fitting all pairs around the central atom. This is commonly sp³d² hybridized. - **For Molecule C:** With three effective pairs involving two sigma bonds and a lone pair, molecule C adopts a trigonal planar structure. It aligns with an sp² hybridization as a result of having a blend of different electron interactions.
Understanding effective pairs allows one to predict the molecular shape and explains the interactions and bond formations within a molecule.