Question

Write down the hybridization of the central atom in cach of the following species: (a) \(\mathrm{SiF}_{4} ;(\mathrm{b}) \mathrm{NH}_{3}\) (c) \(\left[\mathrm{NH}_{4}\right]^{+} ;(\mathrm{d}) \mathrm{BH}_{3} ;(\mathrm{e})\left[\mathrm{CoF}_{6}\right]^{3-} ;(\mathrm{f}) \mathrm{IF}_{3} ;(\mathrm{g}) \mathrm{H}_{2} \mathrm{S}\)

Short answer

(a) sp³, (b) sp³, (c) sp³, (d) sp², (e) d²sp³, (f) sp³d, (g) sp³.

Step-by-step solution

Determine Hybridization for SiF4

Silicon tetrafluoride (0SiF_4) has a central silicon atom bonded to four fluorines. Silicon's valence electron configuration is 3s² 3p². The formation of four bonds indicates the hybridization involves four orbitals (one s and three p), resulting in 0sp^30 hybridization.

Determine Hybridization for NH3

Ammonia (0NH_3) has a nitrogen atom at the center bonded to three hydrogens and one lone pair. Nitrogen's valence electron configuration is 2s² 2p³. It uses three orbitals to bond and one for the lone pair, indicating 0sp^30 hybridization.

Determine Hybridization for [NH4]+

In the ammonium ion (0[NH_4]^+0), nitrogen is surrounded by four hydrogens with no lone pairs. Similar to 0NH_30, this formation requires four orbitals in total, leading to 0sp^30 hybridization.

Determine Hybridization for BH3

Boron hydride (0BH_30) has a central boron atom bonded to three hydrogens. Boron's electron configuration is 2s² 2p¹. It forms three bonds, utilizing three orbitals, leading to 0sp^20 hybridization due to the lack of a lone pair.

Determine Hybridization for [CoF6]3-

Cobalt hexafluoride (0[CoF_6]^{3-}0) involves a cobalt atom with typically empty 3d orbitals. Since it's forming six-coordinate bonds, it uses one s orbital, three p orbitals, and two d orbitals, leading to 0d^2sp^30 hybridization.

Determine Hybridization for IF3

In iodine trifluoride (0IF_30), the iodine is bonded to three fluorines and has two lone pairs. This requires five orbitals total, leading to 0sp^3d0 hybridization to accommodate three bonds and the lone pairs.

Determine Hybridization for H2S

In hydrogen sulfide (0H_2S0), sulfur is bonded to two hydrogens and has two lone pairs. Its electron configuration is 3s² 3p⁴, which requires four orbitals for bonding/lone pairs, resulting in 0sp^30 hybridization.

Key concepts

Hybridization

Hybridization in chemistry is the concept where atomic orbitals mix to form new hybrid orbitals. These new orbitals are suitable for the pairing of electrons to form chemical bonds in molecules. Hybridization helps to explain the structure and the changes in energy levels when atoms bond.
For example:

• SiF₄: Silicon's electron configuration is 3s² 3p². It forms four bonds, suggesting the involvement of one s and three p orbitals, leading to sp³ hybridization.
• NH₃: Here, nitrogen forms three bonds and holds one lone pair, using three p and one s orbital, resulting in sp³ hybridization.
• BH₃: Boron binds with three hydrogens using one s and two p orbitals, indicating sp² hybridization.

Understanding hybridization simplifies the prediction of molecule shapes and angles, making it a key topic in chemical bonding.

Molecular Geometry

The geometry of molecules determines the spatial arrangement of atoms, which can be predicted using hybridization concepts and VSEPR (Valence Shell Electron Pair Repulsion) theory. Different types of hybridizations lead to different molecular geometries.
For instance:

• Tetrahedral Geometry: Observed in compounds like SiF₄ and ext{[NH}_4]^+, with sp³ hybridization forming 109.5° bond angles.
• Trigonal Planar: Seen in BH₃, with sp² hybridization, putting atoms in a plane at 120° angles.
• T-shaped: In IF₃, which involves sp³d hybridization, accommodating five electron domains but resulting in a distorted geometry due to lone pairs.

Comprehending molecular geometry helps in predicting the chemical reactivity and properties of a compound.

Valence Bond Theory

Valence Bond Theory explains how atoms come together to form a molecule by overlapping their atomic orbitals, forming chemical bonds. It underscores the role of valence electrons in bond formation.
This theory articulates:

• Individual atoms retain their orbitals when they bond.
• Overlap of orbitals leads to bond formation, influencing molecular stability.
• Hybridization is a key component highlighting how orbitals adapt and mix in the bonding process.

For example, in NH₃, nitrogen's unhybridized p orbitals overlap with hydrogen's s orbitals to form bonds, each pair of electrons in these overlaps leading to a sigma bond. Valence Bond Theory provides a quantum mechanics-based understanding of the covalent bonding process.

Electron Configuration

Electron configuration is the distribution of electrons in an atom's orbitals, crucial for understanding how bonds form and predicting chemical behavior. Each element has a unique electron configuration, dictating its chemical properties.
Key points include:

• Aufbau Principle: Electrons fill orbitals starting at the lowest energy level.
• Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers in an atom.
• Hund's Rule: Every orbital in a subshell gets single occupancy before any gets double.

Using these rules, one can determine hybridization. Take SiF₄: Silicon’s electron configuration with 3s² 3p² leads to sp³ hybridization, since it requires four hybrid orbitals to bond with four fluorine atoms. Properly understanding electron configurations aids in deciphering molecular structures and reactions.