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Chemical bonding is the fundamental concept that explains how atoms combine to form compounds. It's an attraction between atoms that enables the formation of chemical substances containing two or more atoms. The bond formation releases energy and results in a stable electronic arrangement for the bonded atoms.

The different types of chemical bonds include ionic, covalent, and metallic bonds. Covalent bonding, which Valence Bond Theory (VBT) primarily deals with, involves the sharing of electron pairs between atoms. This sharing allows each atom to gain the electrons it requires to obtain a stable electron configuration reminiscent of the noble gases.

When atoms bond to form molecules, their atomic orbitals, the regions in space where the probability of finding electrons is highest, combine to create molecular orbitals. Molecular orbitals are the conceptual regions within a molecule where shared electrons reside.

A molecular orbital can be bonding, antibonding, or non-bonding, and it can accommodate a maximum of two electrons with opposite spins. The formation of molecular orbitals is central to molecular orbital theory, which describes how electrons are distributed in a molecule and determines the chemical and physical properties of substances.

Orbital overlap is a key concept in covalent bonding as per Valence Bond Theory. For a covalent bond to form between two atoms, their orbitals must overlap sufficiently to allow for the sharing of electrons.

The extent of overlapping influences the strength and characteristics of the bond. A greater overlap results in a stronger bond, which is why the bond length decreases as the bond order increases. According to VBT, the directional nature of orbital overlap is responsible for the shape and geometry of molecules.

A covalent bond is a strong chemical bond formed through the sharing of one or more pairs of electrons between two atoms in a molecule. This sharing allows each atom to fill its outermost electron shell, achieving a more stable arrangement.

Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs. Valence Bond Theory helps explain these bonds by assuming that electron-pair bonds form between atoms through the overlap of their half-filled atomic orbitals.

The hydrogen molecule \(\mathrm{H}_{2}\) serves as the simplest model for explaining covalent bonding. In its diatomic form, two hydrogen atoms share their single electrons, resulting in a covalent bond. According to VBT, the 1s atomic orbitals of each hydrogen atom overlap to create a filled bonding molecular orbital. This overlap accounts for the strong bond holding the hydrogen atoms together in a molecule, illustrating the fundamental principles of chemical bonding.

This bonding results in a stable electron configuration and a reduction in potential energy, which is why \(\mathrm{H}_{2}\) exists as a stable diatomic gas under normal conditions. Understanding the formation of the hydrogen molecule is a stepping stone to comprehending more complex bonding scenarios.