How does each of the three major bonding theories define a single chemical bond?

Dot structures denote a single chemical bond, and curly arrows depict the motion of electrons. Ionic and valence bonds are the two ways that an element's electrons can create a chemical bond (Tro). Ionic compounds contain two positively and negatively charged elements. In order to create a bond, the positively charged atom gives up its electrons, while the negatively charged atom accepts them. Atoms of various elements share electrons equally in valence bonding (Chang). According to this theory, a chemical bond will form through the overlap of atomic orbitals of the participating atoms. A covalent bond form between two atoms through the overlapping of half-filled valence atomic orbitals from individual atoms. Because of overlapping, electrons become localized in the bond regions. Overlapping of atomic orbitals takes different forms. A sigma bond may be formed when atomic orbitals such as s-s, s-p, s-d, p-p, p-d, and d-d overlap (Silberberg).

Molecular orbital theory

In this theory, chemical bonds are formed through the combination of atomic orbitals of the atoms within a molecule. An example is in a Hydrogen molecule (H2), where a single molecular orbital is constructed through the addition of mathematical functions for the two 1s atomic orbitals, which approach each other forming a molecule (Silberberg). Another molecular orbital is created through the subtraction of one of the functions from the other (Petrucci, Herring and Madura).

How does each of the three major bonding theories define a double chemical bond?

Lewis model

Double bonds are represented using double lines between the atoms that have bonded. To form a double bond, all valence electrons involved are added up. The central atom is chosen and a skeletal structure created where atoms are joined using dashes to represent bonds. Electrons are added in pairs starting from the outer electrons until the maximum number of electrons is reached. The remaining electrons are used to form the second bond by either sharing or donor and acceptor atoms (Petrucci, Herring and Madura).

Valence bond theory

Using VBT, a double bond is described using a sigma bond linked to the hybrid orbitals and a pi bond that is linked to unhybridized p orbitals. For instance, in CO2 bonding, C has 4 electrons, where two are 2s, and two are 2p. O has 6 electrons, two 2s, and four 2p. CO2 has a double bond. So as to understand the nature of the bonds, hybridization must be done (Tro). During hybridization of C, the two-2p orbitals remain unhybridized. The hybridized 2s orbitals form sigma bond while the unhybridized 2p orbitals form the pi bond.

Molecular orbital theory

During formation of a molecular with a double bond, the amount orbitals in the molecule is equal to the number of orbitals of atoms combining. When two atoms, each having a single atomic orbital, come together, two molecular orbitals are created. One has small energy than the atomic orbitals (bonding orbital), and the other has higher energy than the atomic orbitals (antibonding orbital). In double bonding, two pairs of electrons are involved in bonding and yields into sigma and pi bonds (Tro).

How does each of the three major bonding theories define a triple chemical bond?

Lewis model

A triple bond is created when 3 electron pairs become shared between the pair of atoms, such as in CO. Valence electrons are determined, and a skeletal structure of the ion or molecule is drawn. These outer shell electrons are arranged around the central atoms, and each atom is connected to the central atom using a pair of electrons or a single bond (Chang). The remaining electrons are distributed as lone pairs on terminal atoms. Electrons are rearranged on the outer atoms forming multiple bonds with the central atom.

Valence bond theory

VBT is used to explain the formation of the triple bond using hybridization. A triple bond contains one sigma bond and two pi bonds. Using ethylene (C2H2) as an example, two C atoms are central and have triple bonds to each other. The arrangement is linear due to sp hybrid orbitals, where only 2 out of 4 electrons of each C are hybridized. The remaining two electrons are on unhybridized p orbital. Thus, there is sp (C) and s (H) overlap and with C atom of other sp creating two sigma bonds besides the two unhybridized p electrons 900 to each other. The coming together of C atoms creates a side overlap of the two electrons on one pi bond, where another pi bond is formed (Petrucci, Herring and Madura).

Molecular orbital theory

Using Hydrogen Fluoride as an example of a compound with a triple bond. MOT explains triple bond in that the two highest energy molecular orbitals are degenerate and are of p-type, having no electron density linked to the H atom, since it is a non-bonding orbital. Thus, the formation of triple bond encompass antibonding, non-binding, and bonding orbitals (Silberberg).

What are the Similarities and differences of Lewis Model, Valence Bond, and Molecular Orbital Theories?

Lewis model helps in the prediction of polyatomic ion and molecular geometries and is based on Lewis structures and repulsion of the charge clouds on the central atoms. Valence bond theory explains the formation of bonds through atomic orbital overlaps and hybrid orbitals such as sp, sp2, sp3, sp3d1, sp3d2 (Petrucci, Herring and Madura). It provides a prediction of molecular geometries, hybrid orbitals, and bond angles. The molecular orbital theory assumes that a molecule is a single entity instead of a collection of hybrid and atomic orbitals. Unlike valence bond theory that takes into consideration atomic orbitals, only molecular orbitals are considered in this theory. Also, electrons are distributed over the entire molecule and not confined to individual atoms. In Lewis model, geometries are inconsistent with s, p, and d orbital shapes. VBT provides an explanation of the discrepancies of orbital shapes through s, p, and d hybridization concepts (Tro). The theory does not give the reason why O2 has a paramagnetic nature. MO theory provides an explanation for the paramagnetic nature of O2.















References

Chang, Raymond. General Chemistry: The Essential Concepts. New York: McGraw-Hill Higher Education, 2002.

Petrucci, Ralph H., et al. General Chemistry: Principles and Modern Applicationss. 10th. Ontario, Canada: Pearson Prentice Hall, 2011.

Silberberg, Martin. Principles of General Chemistry. McGraw-Hill Education: New York, 2012.

Tro, Nivaldo J. Chemistry: A Molecular Approach. London: Pearson, 2016. Online.