Lewis theory single covalent bond

Lewis theory helps us to understand why elemental hydrogen and chlorine exist as diatomic molecules, H2 and CI2. In each case, a pair of electrons is shared between the two atoms. The sharing of a single pair of electrons between bonded atoms produces a single covalent bond. To underscore the importance of electron pairs in the Lewis theory the term bond pair applies to a pair of electrons in a covalent bond, while lone pair applies to electron pairs that are not involved in bonding. Also, in writing Lewis structures it is customary to replace bond pairs with lines (—). These features are shown in the following Lewis structures. 

A stable molecular species has more electrons in bonding orbitals than in antibonding orbitals. For example, if the excess of bonding over antibonding electrons is tzvo, this corresponds to a single covalent bond in Lewis theory. In molecular orbital theory, we say that the bond order is 1. Bond order is one-half the difference between the number (no.) of bonding and antibonding electrons (e ), that is. 

Valence bond (VB) theory, which helps explain bonding and structure in main-group compounds (Section 11.1), is also used to describe bonding in complex ions. In the formation of a complex ion, the filled ligand orbital overlaps the empty metal-ion orbital. The ligand (Lewis base) donates the electron pair, and the metal ion (Lewis acid) accepts it to form one of the covalent bonds of the complex ion (Lewis adduct) (Section 18.8). Such a bond, in which one atom in the bond contributes both electrons, is called a coordinate covalent bond, although, once formed, it is identical to any covalent single bond. Recall that the... 

Molecular orbital (MO) theory combines the tendency of atoms to fill their octets by sharing electrons (the Lewis model) with their wavelike properties—assigning electrons to a volume of space called an orbital. According to MO theory, covalent bonds result from the combination of atomic orbitals to form molecular orbitals— orbitals that belong to the whole molecule rather than to a single atom. Like an atomic orbital that describes the volume of space around the nucleus of an atom where an electron is likely to be found, a molecular orbital describes the volume of space around a molecule where an electron is likely to be found. Like atomic orbitals, molecular orbitals have specific sizes, shapes, and energies. 

Two types of bonds, namely, ionic and nonionic (covalent) bonds, were recognized early on in the formulation of the electronic theory of valency. Lewis made clear distinctions between ionic and covalent bonds. The first are formed by transfer of electrons and production of separate charged ions (as explained by Kossel, 1916), the second, according to Lewis, by sharing of electrons in pairs, a single bond consisting of one shared pair, a double bond of two, and a triple bond of three. 

The Lewis theory of bonding describes a covalent bond as the sharing of a pair of electrons, but this does not necessarily mean that each atom contributes an electron to the bond. A covalent bond in which a single atom contributes both of the electrons to a shared pair is called a coordinate covalent bond. 

H2 This molecule has two electrons, both in the orbital. The bond order is (2 — 0)/2 = 1. With Lewis theory and the valence bond method, we describe the bond in H2 as single covalent. 

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