The VSEPR model is usually a satisfactory method for predicting molecular geometries.To understand bonding and electronic structure, however, you must look to quantum mechanics. We will consider two theories stemming from quantum mechanics:
valence bond theory and molecular orbital theory. Both use the methods of quantum mechanics but make different simplifying assumptions. In this lecture today, we will look in a qualitative way at the basic ideas involved in valence bond theory, an approximate theory to explain the electron pair or covalent bond by quantum mechanics.
According to valence bond theory, a bond forms between two atoms when the following
conditions are met:
1. An orbital on one atom comes to occupy a portion of the same region of space
as an orbital on the other atom. The two orbitals are said to overlap.
2. The total number of electrons in both orbitals is no more than two.
As the orbital of one atom overlaps the orbital of another, the electrons in the orbitals begin to move about both atoms. Because the electrons are attracted to both nuclei at once, they pull the atoms together. Strength of bonding depends on the amount of overlap; the greater the overlap, the greater the bond strength. The two orbitals cannot contain more than two electrons because a given region of space can hold only two electrons with opposite spin.
For example, consider the formation of the H2 molecule from atoms. Each atom has the electron configuration 1s1. As the H atoms approach each other, their 1s orbitals begin to overlap and a covalent bond forms (See Figure Below)
. Valence bond theory also explains why two He atoms (each with electron configuration 1s2) do not bond. Suppose two He atoms approach one another and their 1s orbitals begin to overlap. Each orbital is doubly occupied, so the sharing of electrons between atoms would place the four valence electrons from the two atoms in the same region. This, of course, could not happen. As the orbitals begin to overlap, each electron pair strongly repels the other. The atoms come together, then fly apart.
Because the strength of bonding depends on orbital overlap, orbitals other than s orbitals bond only in given directions. Orbitals bond in the directions in which they protrude or point, to obtain maximum overlap. Consider the bonding between a hydrogen atom and a chlorine atom to give the HCl molecule. A chlorine atom has the electron configuration [Ne]3s23p5. Of the orbitals in the valence shell of the chlorine atom, three are doubly occupied by electrons and one (a 3p orbital) is singly occupied. The bonding of the hydrogen atom has to occur with the singly occupied 3p orbital of chlorine. For the strongest bonding to occur, the maximum overlap of orbitals is required. The 1s orbital of hydrogen must overlap along the axis of the singly occupied 3p orbital of chlorine.
From what has been said, you might expect the number of bonds formed by a given atom to equal the number of unpaired electrons in its valence shell. Chlorine, whose orbital diagram is has one unpaired electron and forms one bond. Oxygen, whose orbital diagram is has two unpaired electrons and forms two bonds, as in H2O.
1s 2s 2p
1s 2s 2p
Hybrid orbitals are orbitals used to describe bonding that are obtained by taking combinations
of atomic orbitals of the isolated atoms. In this case, a set of hybrid orbitals is constructed from one s orbital and three p orbitals, so they are called sp3 hybrid orbitals. Calculations from theory show that each sp3 hybrid orbital has a large lobe pointing in one direction and a small lobe pointing in the opposite direction. The four sp3 hybrid orbitals point in tetrahedral directions.
The C-H bonds in methane, CH4, are described by valence bond theory as the
overlapping of each sp3 hybrid orbital of the carbon atom with 1s orbitals of hydrogen
atoms .Thus, the bonds are arranged tetrahedrally, which is predicted by the VSEPR model.
Hybrid orbitals can be formed from various numbers of atomic orbitals.
The number of hybrid orbitals formed always equals the number of atomic orbitals used. For
example, if you combine an s orbital and two p orbitals to get a set of equivalent orbitals, you get three hybrid orbitals (called sp2 hybrid orbitals). A set of hybrid orbitals always has definite directional characteristics. Here, all three sp2 hybrid orbitals lie in a plane and are directed at 120 angles to one another; that is, they have a trigonal planar arrangement
Recommended lecture: Effect of Polarity on Molecular Properties
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