Question

What is the hybridization of the central atom in each of the following molecules? a. \(\mathrm{SF}_{6}\) b. \(\mathrm{ClF}_{3}\) c. \(\mathrm{GeCl}_{4}\) d. \(\mathrm{XeF}_{4}\)

Short answer

The hybridization of the central atom for each molecule is as follows: a. \(\mathrm{SF}_{6}\): \(\mathrm{sp}^{3}\mathrm{d}^{2}\) b. \(\mathrm{ClF}_{3}\): \(\mathrm{sp}^{3}\mathrm{d}\) c. \(\mathrm{GeCl}_{4}\): \(\mathrm{sp}^{3}\) d. \(\mathrm{XeF}_{4}\): \(\mathrm{sp}^{3}\mathrm{d}^{2}\)

Step-by-step solution

a. \(\mathrm{SF}_{6}\)#

To find the central atom's hybridization in \(\mathrm{SF}_{6}\), we first need to count the electron groups around the central sulfur (S) atom. Since S is bonded to six fluorine (F) atoms, there are six electron groups. Hence, the hybridization is \(\mathrm{sp}^{3}\mathrm{d}^{2}\), since it satisfies the requirement of six orbitals (1 s, 3 p, and 2 d orbitals).

b. \(\mathrm{ClF}_{3}\)#

For the central atom, chlorine (Cl), in \(\mathrm{ClF}_{3}\), we count the electron groups around it. Cl is bonded to three fluorine (F) atoms. Additionally, Cl has two lone pairs, summing up to a total of five electron groups. The hybridization is \(\mathrm{sp}^{3}\mathrm{d}\), which accounts for the required five orbitals (1 s, 3 p, and 1 d orbitals).

c. \(\mathrm{GeCl}_{4}\)#

To find the central atom's hybridization in \(\mathrm{GeCl}_{4}\), we need to count the electron groups around the central germanium (Ge) atom. Germanium is bonded to four chlorine (Cl) atoms, so there are four electron groups. Therefore, the hybridization is \(\mathrm{sp}^{3}\), which accounts for the four orbitals needed (1 s and 3 p orbitals).

d. \(\mathrm{XeF}_{4}\)#

For the xenon (Xe) central atom in \(\mathrm{XeF}_{4}\), we need to count the electron groups around it. Xe is bonded to four fluorine (F) atoms and has two lone pairs, so there are a total of six electron groups. The hybridization is \(\mathrm{sp}^{3}\mathrm{d}^{2}\), which accounts for the necessary six orbitals (1 s, 3 p, and 2 d orbitals).

Key concepts

Electron Groups

Electron groups are essential to understanding molecular geometry and hybridization. They consist of all atoms and lone pairs surrounding a central atom.
In valence shell electron pair repulsion (VSEPR) theory, electron groups arrange themselves to minimize repulsion. This arrangement is vital for predicting the shape of molecules.

• Bonding pairs: These are electron groups involved in chemical bonds.
• Lone pairs: Electrons not involved in bonding, they are an electron group occupying space around the central atom.

Identifying the number of electron groups helps in determining hybridization states. For example, in \(SF_6\), there are six electron groups, leading to an \(sp^3d^2\) hybridization, ensuring maximum spatial separation and the corresponding molecular shape.

Molecular Geometry

Molecular geometry describes the spatial arrangement of atoms within a molecule. It determines the molecule's shape and is heavily influenced by the number of electron groups.
VSEPR theory helps us predict this shape based on both bonding pairs and lone pairs.

• Linear geometry occurs when there are two electron groups.
• Tetrahedral results from four electron groups.
• Trigonal bipyramidal is associated with five electron groups.
• Octahedral comes from six electron groups, like in \(SF_6\).

Understanding molecular geometry is crucial for grasping how molecules interact with each other, affecting their physical and chemical properties.

Orbital Overlap

Orbital overlap refers to the way atomic orbitals mix to form new hybrid orbitals, which will hold the electron pairs that form chemical bonds.
This concept is where valence bond theory plays a role. When orbitals overlap, they form a strong bond in the molecule.

• Sigma (σ) bonds are formed by head-on overlapping, providing a strong bond along the axis between two nuclei.
• Pi (π) bonds are formed by side-on overlapping, which occurs above and below the axis of the bonding atoms.

The effectiveness of the overlap affects the strength of the chemical bond. For instance, in \(GeCl_4\), the overlapping of s and p orbitals forms strong covalent bonds in a tetrahedral shape.

Chemical Bonding

Chemical bonding involves the interaction of atoms to achieve stable electron configurations. The main types of chemical bonds include ionic, covalent, and metallic bonds.
Hybridization is a key concept to explain covalent bonding, as it describes how different orbitals combine to form equivalent hybrid orbitals.

• In \(ClF_3\), chlorine forms covalent bonds with fluorine, sharing electron pairs.
• Hybridization helps explain how these bonds are formed, with lone pairs also playing a role in determining the bond angles and molecular geometry.
• In \(XeF_4\), even noble gases engage in chemical bonding due to hybridization.

Understanding chemical bonding gives insight into molecular structure, reactivity, and properties. It is foundational in chemistry for explaining how molecules are constructed and behave.