A covalent bond develops when the orbitals of two atoms overlap and a pair of electrons occupy the overlap area, according to the basic premise of VB theory.

The overlap of the two orbitals indicates that their wave functions are in phase (constructive interference; as referenced to the image attached), and therefore the amplitude between the nuclei rises. This idea underpins the major themes of VB theory:

The electron pair is opposing spins.

The space produced by the overlapping orbitals has a maximum capacity for two electrons with opposing (paired) spins, as required by the Pauli exclusion principle.

A molecule of H2 develops in the simplest scenario when the 1s orbitals of two H atoms overlap and the electrons, with their spins coupled, spend more time in the overlap area (up and down arrows in the attached picture).

Bonding orbitals with the greatest overlap. Connection strength is determined by the attraction of nuclei to shared electrons; hence, the higher the orbital overlap, the closer the nuclei are to the electrons, and the stronger the bond.

The extent of overlap is determined by the orbital shape and orientation.

Atomic orbital hybridization To explain bonding in diatomic compounds such as HF and F2, we imagine a direct overlap of s and/or p orbitals of isolated atoms.

But how can the shapes and orientations of C and H atomic orbitals account for the structure of a molecule like methane? A-C atom ([He] 2s 2 2p2) contains two valence electrons in the spherical 2s orbital and one in two of the three mutually perpendicular 2p orbitals.

Two CH bonds with a 90° HCH bond angle would occur if the two half-filled p orbitals overlapped the 1s orbitals of two H atoms.

Methane, on the other hand, has the formula CH4, not CH2, and its bond angles are 109.5°. Linus Pauling proposed a theory to explain such phenomena.

Hybrid orbitals have certain characteristics. Here are some key facts to remember about the hybrid orbitals that develop during bonding:

The number of hybrid orbitals created matches the number of atomic orbitals mixed, and the kind of hybrid orbitals formed vary with the types of atomic orbitals combined.

A hybrid orbitals shape and orientation optimize its overlap with the orbital of the other atom in the bond.

It's helpful to conceive of hybridization as a process in which atomic orbitals mix, hybrid orbitals develop, they overlap other orbitals, and electrons with opposing spins join the overlap zone, establishing stable bonds.

In reality, hybridization is a mathematical notion that helps us understand the chemical world.

The VSEPR and VB theories are used to explain observed molecule shapes. However, in certain situations, the ideas may be inconsistent with other data.

Large nonmetal hydrides are not suitable for hybridization. Consider H2S's Lewis structure and bond angle:

According to VSEPR theory, the four-electron groups surrounding H2S point to the corners of a tetrahedron, and the two lone pairs compress the HSH bond angle below the ideal 109.5°.

According to VB theory, the S atom's 3s and 3p orbitals combine to create four sp3 hybrids, two of which are filled with lone pairs while the other two overlap 1s orbitals of two H atoms and are filled with bonding pairs.

Shapes with enlarged valence shells are less significant for d-Orbital hybridization. The creation of sp3 d and sp3 d2 hybrid orbitals with enlarged valence shells is proposed by the VB theory.

However, new quantum-mechanical simulations have revealed that d orbitals have such high energy that they do not successfully hybridize with the considerably more stable s and p orbitals of a given n number.

Thus, it now appears that SF6 does not bind using sp3 d2 hybrids; instead, some have hypothesized that SF6 is best to the stable when the bonding orbitals of the central S employ a combination of sp hybrid orbitals and unhybridized 3p orbitals.

Others favor molecular orbitals or even ionic structures as explanations. We must remember that molecules do not have to obey our models; rather, we must understand the limitations of our models and change them in light of new data.

Because VB theory successfully explains the molecular geometries of molecules with expanded valence shells, we will continue to use the traditional approach of including d-orbital hybridization for molecules with expanded valence shells in this text for simplicity, while acknowledging its limitations.

End-to-end overlap of atomic orbitals produces a connection, allowing unrestricted rotation of the molecule's linked components. A multiple bond is made up of a bond and either one (double bond) or two bonds (triple bond).

Multiple bonds contain more electron density between their nuclei than single bonds, resulting in higher bond energies. Rotation in a bond is limited by the side-to-side overlap of orbitals.

Molecular orbital (MO) theory considers a molecule to be a collection of nuclei with MOs distributed throughout the structure.

Atomic orbitals with equivalent energies can be combined or removed to produce bonding or antibonding MOs.

Bonding MOs, whether or, have most of the electron density between the nuclei and are lower in energy than the AOs that combine to create them; antibonding MOs have most of the electron density outside the nuclei and are thus greater in energy.

MOs are filled in descending order of energy by paired electrons with opposing spins. MO graphs depict energy levels as well as orbital occupancy. Period 2 homonuclear diatomic molecular diagrams

Several features of MO theory, including MOThe charge filling, energy-level diagrams, electron configuration, and bond order, are related to previous concepts: Electrons are injected into MOs. Electrons enter MOs in the same way as they do AOs:

MOs are filled in ascending energy order (Aufbau principle). An MO can only contain two electrons with opposing spins (Pauli exclusion principle).

Before any of the orbitals of equal energy are filled, they are half-filled, with electron spins parallel (Hund's rule).