10.1 - Molecular Geometry

• This method is based on the idea that electron pairs in an atom's valence shell reject one another.

  • The valence shell is an atom's outermost electron-occupied shell, and it contains the electrons engaged in bonding.

  • The valence-shell electron-pair repulsion (VSEPR) model is a method of studying molecular geometry

    • It accounts for the geometric arrangements of electron pairs around a core atom in terms of electron pair electrostatic repulsion.

10.2 - Dipole Moments

• The dipole moment, which is the product of the charge Q and the distance r between the charges, is a quantitative measure of a bond's polarity: µ = Q × r

• Polar molecules are diatomic compounds with dipole moments that comprise atoms of various elements (for example, HCl, CO, and NO).

  • Diatomic molecules are made up of the same element's atoms.

  • Because they lack dipole moments, nonpolar molecules are nonpolar.

10.3 - Valence Bond Theory

• Covalent bond formation and molecule electronic structure are described using two quantum mechanical theories.

  • The valence bond (VB) theory states that the electrons in a molecule inhibit the individual atoms' atomic orbitals.

  • It allows us to keep track of specific atoms involved in the bonding process.

  • The molecular orbital (MO) theory, on the other hand, posits that molecular orbitals are formed from atomic orbitals.

10.4 Hybridization of Atomic Orbitals

• VB theory uses hypothetical hybrid orbitals to explain methane bonding.

  • These are atomic orbitals formed when two or more nonequivalent orbitals of the same atom join in preparation for covalent bond formation.

  • The term "hybridization" refers to the process of combining atomic orbitals in an atom (typically a central atom) to produce a set of hybrid orbitals.

10.5 - Hybridization in Molecules Containing Double and Triple Bonds

• In Figure 10.16(a), each C atom forms three sigmas () bonds, which are covalent bonds created by orbitals overlapping end-to-end with the electron density centered between the nuclei of the bonding atoms.

• A pi bond is a covalent link produced by sideways overlapping orbitals with electron density concentrated above and below the plane of the connecting atoms' nucleus.

10.6 - Molecular Orbital Theory

• Another quantum mechanical technique known as molecular orbital (MO) theory can occasionally explain the magnetic and other properties of molecules.

  • The molecular orbital theory defines covalent bonds in terms of molecular orbitals, which are associated with the entire molecule and result from the interaction of the atomic orbitals of the bonding atoms.

• The overlap of the 1s orbitals of two hydrogen atoms, according to MO theory, results in the development of two molecular orbitals: one bonding and one antibonding.

  • The energy and stability of a bonding molecular orbital are lower than those of the atomic orbitals from which it was generated.

  • Antibonding molecular orbitals have more energy and are less stable than the atomic orbitals from which they were created.

• The electron density in a sigma molecular orbital is focused symmetrically along a line connecting the two nuclei of the bonding atoms.

• The electron density is concentrated above and below a line connecting the two nuclei of the bonding atoms in a pi molecular orbital.

10.7 - Molecular Orbital Configurations

• The total number of atomic orbitals formed is always equal to the total number of molecular orbitals formed.

• The related antibonding molecular orbital is less stable than the bonding molecular orbital.

• The filling of molecular orbitals begins at low energy and progresses to higher energies.

• The number of electrons in bonding molecular orbitals is always greater than that in antibonding molecular orbitals in a stable molecule.

• According to the Pauli exclusion principle, each molecular orbital can hold up to two electrons with opposite spins, just as an atomic orbital.

• Hund's rule predicts the most stable arrangement when electrons are added to molecular orbitals of the same energy; that is, electrons enter these molecular orbitals with parallel spins.

• The sum of all the electrons on the bonding atoms equals the number of electrons in the molecular orbitals.

10.8 - Delocalized Molecular Orbitals

• In contrast to the pi bonding molecular orbitals in ethylene, those in benzene create delocalized molecular orbitals, which are not limited between two nearby bonding atoms but instead span three or more.