Resonance
The phenomenon in which a molecule or ion cannot be represented by a single structure but represented by more than one structure to explain its properties is called resonance. More than one such structure of the same molecule or Ion is called a resonating structure. The real structure of the molecule or ion is a hybrid of these structures called a resonance hybrid. It is represented by a double-headed arrow (↔).
Examples:
1. Resonance in ozone
2. Resonance in NO2–
Valence shell electron pair repulsion (VSEPR) theory
To predict the shape of the covalent molecule, Gillespie and Nyholm developed a theory known as the valence shell electron pair repulsion theory. The main postulates of this theory are given below:
1. The shape of covalent molecules depends upon the number and nature of electron pairs surrounding the central atom. The electron pairs are arranged as far apart as possible so that there is minimum repulsion between them due to which stability is increased.
2. If the central atom of a molecule is surrounded only by a bond pair of electrons, the shape of the molecule is regular.
| Number of bond pairs | Shape |
|---|---|
| 2 | Liner |
| 3 | Trigonal Planar |
| 4 | Tetrahedral |
| 5 | Trigonal bipyramidal |
| 6 | Octahedral |
3. If the central atom of a molecule is surrounded by both a bond pair and a lone pair of electrons, the molecule does not have a regular shape. It is due to the presence of a lone pair of electrons. The repulsive interaction between the electron pair follows the following order: lp-lp>lp-bp>bp-bp.
4. The electronegativity of atoms in a molecule also affects the bond angle of the molecule.
- As the electronegativity of the central atom increases, the bond angle increases.
- As the electronegativity of atoms bonded to the central atom increases, the bond angle decreases.
5. This theory considers a multiple bond as a single-bond pair of electrons.
6. Repulsion between electron pairs in completely filled shell is larger than in incompletely filled shell.
Examples:
i. BeF2 molecule
The central atom Be of the BeF2 molecule is surrounded by two bond pairs of electrons. To minimise the repulsion between electron pairs, they must be arranged linearly at 180°. So the molecule has a linear shape.
\underset {Linear\ shape}{F - Be - F}ii. BF3 molecule
The central atom B of the BF3 molecule is surrounded by three bond pairs of electrons. To minimise the repulsion between electron pairs, they must be directed towards three corners of an equilateral triangle with a bond angle of 120°. So the molecule has a trigonal planar shape.
iii. CH4 molecule
The central atom C of the CH4 molecule is surrounded by four bond pairs of electrons. To minimise the repulsion between electron pairs, they must be directed towards four corners of a regular tetrahedron with a bond angle of 109.5°. So the molecule has a tetrahedral shape.
iv. PCl5 molecule
The central atom P of the PCl5 molecule is surrounded by five bond pairs of an electron. To minimise the repulsion between them, three of them lie in an equilateral plane with a 120° angle. The remaining two electron pairs are arranged perpendicularly opposite to this plane. So the shape of the molecule is trigonal bipyramidal.
v. SF6 molecule
The central atom of the molecule is surrounded by six bond pairs of an electron. To minimise the repulsion between them, four of them lie in an equatorial plane with an angle of 90°. The remaining two Bond pairs are arranged perpendicularly and opposite to the equatorial plane. So the shape of the molecule is octahedral.
vi. NH3 molecule
The central atom N of the NH3 molecule is surrounded by 4 electron pairs. To minimise the repulsion between them, they should be arranged in a tetrahedral manner. So the expected shape of the molecule is tetrahedral. But, due to the presence of one lone pair of an electron, there is distortion in the shape because the lone pair exerts greater repulsion to the bond pair due to which contraction in bond angle takes place. So the actual shape of the molecule is pyramidal with an angle of 107.5°.
vii. H2O molecule
The central atom O of the H2O molecule is surrounded by 4 electron pairs. To minimise the repulsion between them, they should be arranged in a tetrahedral manner. So the expected shape of the molecule is tetrahedral. But, due to the presence of two lone pairs of electrons, there is distortion in the shape because the lone pair exerts greater repulsion to the bond pair due to which contraction in bond angle takes place. So the actual shape of the molecule is angular with an angle of 105.5°.
Valence Bond Theory
To explain the stability of the covalent molecule, a new method to form a covalent bond was proposed by Hitler and London. The following are the postulates of this theory:
- Covalent bonds are formed due to the overlapping of half-filled atomic orbitals are present in the valence shell of participating atoms.
- For overlapping, the half-filled orbitals must have electrons of opposite spin.
- After overlapping, the pairing of electrons takes place which decreases energy and increases the stability of the system.
- Due to overlapping, the electron density between the two nuclei of the bonded atom is increased which causes the formation of a covalent bond.
- The strength of the covalent bond depends upon the extent of overlapping. Greater the extent of overlapping, stronger is the bond and vice versa.
According to this theory, the covalency of an element is equal to the number of half-filled orbitals present in the valence shell of an atom.
Types of covalent bond
Depending upon the nature of overlapping, there are two types of covalent bonds.
1. Sigma (σ) bond:
A covalent bond formed due to head-on overlapping of half-filled atomic orbitals is called a sigma bond. It is formed due to the following types of overlapping:
i. S-S overlap
ii. S-P overlap
iii. P-P overlap
2. Pi (π) bond
A covalent bond formed due to the sidewise overlapping of half-filled atomic orbitals is called a Pi bond.
Differences between sigma (σ) and pi (π) bonds
| σ – bond | π – bond |
|---|---|
| It is formed by end to end, head-on or axial overlapping of half-filled orbitals. | It is formed by sidewise or lateral overlapping of half filled orbitals. |
| Overlapping of orbitals takes place along the internuclear axis. | Overlapping of orbitals takes place perpendicular to the internuclear axis. |
| The extent of the overlapping is large. So the bond formed is strong. | The extent of the overlapping is small. So the bond formed is weak. |
| It is formed by the overlapping of s-s, s-p or p-p orbitals. | It is formed by the overlapping of p-p orbitals. |
| The molecular orbital is continuous containing only one charge cloud. | The molecular orbital is discontinuous containing two charge clouds. |
| There is a free rotation of atoms around the sigma bond. | There is no free rotation of atoms around the pi bond. |
| It can exist alone. | It does not exist alone. |
Q. How many sigma and pi bonds are present in a. methane b. ethane c. ethene d. ethyne?
Some Important Questions
i. Define resonance. Draw the resonating structure of SO3, CO3—, SO4—.
ii. What are the postulates of VSEPR theory?
iii. BCl3 is trigonal planar but PCl3 is trigonal pyramidal. Why?
iv. NH3 is trigonal pyramidal but NH4 is tetrahedral. Why?
v. What are the differences between the σ and π bonds?
Questions & discussion
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