Valence Bond Theory
This theory was proposed by Heitler and London (1927) and further developed by Pauling and others.
To understand this theory, let us assume that two hydrogen atoms A and B are approaching each other. Their nuclei are NA and NB and electrons are eA and eB. When two atoms come closer to each other, new attractive and repulsive forces begin to work.
Attractive forces arise between:
- Nucleus of one atom and its own electron, i.e. NA-eA and NB-eB
- Nucleus of one atom and electron of other atoms, i.e. NA-eB and NB-eA
Repulsive forces arise between:
- Electrons of two atoms, i.e. eA-eB
- Nuclei of two atoms, i.e. NA-NB
Experiments have shown that magnitude of new attractive force is greater than that of new repulsive force. So, two atoms come closer to each other and potential energy decreases. Finally, the net force of attraction balances the force of repulsion and system acquires minimum energy. At this stage two hydrogen atoms are said to be bonded together to form a stable molecule.
Orbital Overlap Concept: When minimum energy state is achieved, atomic orbitals of two hydrogen atoms undergo partial interpenetration or merger. This partial merging of orbitals is called overalapping of orbitals. The overlapping of orbitals results in pairing of electrons. The extent of overlap decides the strength of a covalent bond. Normally, greater the overlap stronger is the bond. Thus, pairing of electrons (present in the valence shell having opposite spins) results in the formation of covalent bond between two atoms.
Overlapping of Atomic Orbitals: Depending upon the sign (phase) and direction of orientation of amplitude of orbital wave function in space, the overlapping of atomic orbitals may be positive, negative or zero.
- When phase and orientation of two merging orbitals are same, the overlap is called postitive overlap, or in phase overlap.
- When phases are different but orientation is same, it is a negative or out of phase overlap.
- When phases are same but orientation is different, it is a zero overlap or out of phase overlap due to different orientation.
Simple atomic orbital overlap does not account for the directional characteristics of bonds in polyatomic molecules, e.g. CH4, NH3 and H2O
Let us take example of CH4 to understand this.
Electronic configuration of Carbon in ground state: [He] 2s22p2
Electrons configuration of C in excited state: [H] 2s1 2px1 2py1 2pz1
Now, the four atomic orbitals of carbon can overlap with 1s orbitals of four H atoms. This will result in formation of four C-H bonds. While the three p orbitals of C are at 90° to one another, the HCH angle for these will also be 90°. This means that three C-H bonds will be oriented at 90° to one another. The 2s orbital of C and 1s orbital of H are spherically symmetrical and so they can overlap in any direction. Hence, the direction of fourth C-H bond cannot be ascertained. But this description does not fit with the tetrahedral HCH angles of 109.5°.
Types of Overlapping and Nature of Covalent Bond
Depending on the type of overlapping, there are two types of covalent bond: (a) Sigma (σ) bond and(b) pi (π) bond.
(a) Sigma (σ) Bond: This type of bond is formed by end to end (head-on) overlap of bonding orbitals along the internuclear axis. This can be formed by any of the following types of combinations:
- s-s Overlap: Two half-filled s-orbitals overlap.
- s-p Overlap: Half filled s-orbitals of one atom overlap with half filled p-orbital of another aotm.
- p-p Overlap: Half filled p-orbitals overlap.
(b) pi (π) Bond: In this type of overlap, axes of merging orbitals remain parallel to each other and perpendicular to internuclear axis. The orbitals formed due to sidewise overlapping consists two saucer type charged clouds above and below the plane of the participating atoms.
Strength of Sigma and pi bonds: Sigma bond is stronger than pi bond. In case of formation of multiple bonds between two atoms of a molecule, pi bond is formed in addition to sigma bond.
The process of intermixing of orbitals of slightly different energies resulting in the formation of new set of orbitals of equivalent energies and shape is called hybridization. The concept of hybridization was proposed by Pauling and we can explain the shapes of polyatomic molecules on this concept. For example; when one 2s and three 2p orbitals of carbon hybridise, four new sp3 hybrid orbitals are formed.
Key Features of Hybridization:
Number of hybrid orbitals is equal to the number of atomic orbitals that get hybridized.
- Hybrid orbitals are always equivalent in energy and shape.
- Hybrid orbitals are more effective in forming stable bonds than pure orbitals.
- Hybrid orbitals are directed in space in certain preferred direction to have minimum repulsion between electron pairs and thus attain stable arrangement.
Important Conditions for Hybridization
- The orbitals present in the valence shell are hybridized.
- The orbitals undergoing hybridization should have almost equal energy.
- Promotion of electron is not a necessary condition prior to hybridization.
- It is not necessary that only half-filled orbitals participate in hybridization. Sometimes, even filled orbitals of valence shell take part in hybridization.
Types of Hybridization
(a) sp Hybridization: It involves merging of one s and one p orbital resulting in the formation of two equivalent sp hybrid orbitals. Suitable orbitals for sp hybridization are s and pz. Each sp hybrid orbital has 50% s-character and 50% p-character. A molecule in which the central atom is sp-hybridized and is directly linked to two other atoms, the central atom possesses linear geometry.
This is also called diagonal hybridization. The two sp hybrids point in the opposite directions along z-axis. They have projecting positive lobes and very small negative lobes, which provide more effective overlapping resulting in the formation of stronger bonds.
Example of sp hybridization: BeCl2
Ground state electronic configuration of Be: 1s2 2s2
Excited state electronic configuration of Be: 1s2 2s1 2p1
Thus, two sp hybridized orbitals are formed by one 2s and one 2p orbitals. These two sp orbitals are oriented in opposite directions forming an angle of 180°. Each of the sp hybridized ortbitals overlaps with the 2p orbital of chlorine axially and form two Be-Cl sigma bonds.
(b) sp2 Hybridization: It involves merger of one s and two p orbitals resulting in the formation of three equivalent sp2 hybridized orbitals. Molecule of BCL3 gives an example of sp2 hybridization.
Ground state electronic configuration of B: 1s2 2s2 2p1
Excited state electronic configuration of B: 1s2 2s1 2px1 2py1 2pz1
It results in the formation of three sp2 hybridized orbitals which are oriented in a trigonal planar arrangement and overlap with 2p orbitals of chlorine to form three B-Cl bonds.
(c) sp3 Hybridization: It involves merger of one s and three p orbitals. The four sp3 orbitals thus formed are directed towards the four corners of a tetrahedron at an angle of 109.5°. CH4 gives a good example of this type of hybridization.
Hybridization of Elements involving d Orbitals
The energy of 3d orbitals is comparable to that of 3s and 3p orbitals. The energy of 3d orbitals is also comparable to that of 4s and 4p orbitals. So, hybridization involving either 3s, 3p and 3d or 3d, 4s and 4p is possible. But as the energy difference between energies of 3p and 4s orbitals is significant, hybridization involving 3p, 3d and 4s orbitals is not possible. Following table shows important hybridization schemes involving s, p and d orbitals.
|Shape of molecules/ions||Hybridization Type||Atomic Orbitals||Examples|
|Square Planar||dsp2||d + s + p(2)||[Nic(CN)4]2-, [Pt(Cl)4]2-|
|Trigonal Bipyramidal||sp3d||s + p(3) + d||PF5, PCl5|
|Square Pyramidal||sp3d2||s + p(3) + d(2)||BrF5|
|Octahedral||sp3d3, d2sp3||s + p(3) + d(2), d(2) + s + p(3)||SF6, [CrF6]3-, [Co(NH3)6]3+|