NIOS Lesson 4 - CHEMICAL BONDING

In lesson 2, you have learnt about the structure of atom while in the lesson 3, you studied about the classification of elements and periodicity in properties.

You know that molecules are obtained by the combination of two or more than two atoms of the same or different elements. In this lesson you will study

􀁺 Why do atoms combine?

􀁺 What are the different ways in which the atoms can combine? and

􀁺 What are the shapes of different molecules?

The answers to these questions are of fundamental importance to the study of chemistry, as you would discover while studying the later parts of this course.

OBJECTIVES

After reading this lesson you will be able to

􀁺 explain the formation of bond in terms of potential energy diagram and octet rule;

􀁺 list different types of bonds;

􀁺 define ionic bond and cite some examples;

􀁺 write Lewis structures of some simple molecules;

􀁺 list the characteristics of ionic compounds;

􀁺 explain Born Haber Cycle;

􀁺 define covalent bond and cite some examples;

􀁺 list the characteristics of covalent compounds;

􀁺 state valence shell electron pair repulsion (VSEPR) theory;

􀁺 explain bond polarity and dipole moment; Chemical Bonding

􀁺 explain bond parameters;

􀁺 predict the geometry of molecules with the help of VSEPR theory;

􀁺 explain the hybridisation of atomic orbitals involving s, p and d orbitals and illustrate with examples;

􀁺 tabulate the geometry of some molecules showing sp, sp2, sp3, dsp2, and dsp3 hybridisation;

􀁺 explain the formation of σ and π bonds in CH4 , C2H4 and C2H2 ;

􀁺 explain resonance;

􀁺 explain molecular orbital theory;

􀁺 write the molecular orbital configuration of H2, N2, O2 and F2 molecules;

􀁺 define bond length and bond order and relate them and

􀁺 explain hydrogen bonding with the help of examples.

4.1 VALENCE ELECTRONS

4.2 WHAT IS A CHEMICAL BOND?

4.3 IONIC OR ELECTOVALENT BOND

4.3.1 Energetics of Ionic Compound Formation

Born Harber Cycle

4.3.2 Characteristic Properties of Ionic Compounds

4.4 COVALENT BOND

4.4.1 Lewis Structure

4.4.2 Coordinate Covalent Bond

4.4.3 Characteristic properties of Covalent Compounds

4.4.4 Polar Covalent Bond

4.4.5 Bond Polarity and Dipole Moment

Dipole Moment

4.4.6 Covalent Character of Ionic Bond

4.4.7 Covalent Bond Parameters

4.5 HYDROGEN BONDING

4.6 VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY

4.7 MODERN THEORIES OF CHEMICAL BONDING

4.7.1 Valence Bond Theory

4.7.1.1 Hybridisation

4.7.1.2 Hybridisation and Multiple Bonds

4.7.1.3 Resonance

4.7.2 Molecular Orbital Theory

4.7.2.1 Molecular Orbital Bonding in Diatomic Molecules of Second Period

4.7.2.2 MO Electronic Configuration and Properties of a Molecule

WHAT YOU HAVE LEARNT

􀁺 A chemical bond may be visualised as an effect that leads to the decrease in the energy of the combination of two atoms when they come closer.

􀁺 The atoms combine in such a way so as to attain stable electronic configuration of noble gases.

􀁺 According to Kossel, transfer of an electron from one atom to the other achieves the stable configuration. This leads to formation of ions, which are held together by electrostatic interactions called ionic bond.

􀁺 According to Lewis, the stable configuration is achieved by sharing of electron Chemical Bonding pairs between the bonding atoms. This leads to the formation of a covalent bond.

􀁺 Bonding in simple molecules can be conveniently represented in terms of Lewis electron-dot structures.

􀁺 In some covalently bound atoms the shared pair of electron is more towards the atom with greater electronegativity and leads to partial ionic character in the molecule.

􀁺 Valence shell electron pair repulsion (VSEPR) theory is very helpful in predicting the shapes of simple molecules. It is based on the interactions between the electron pairs around the central atom in the molecule.

􀁺 Valence bond theory (VBT) and Molecular orbital theory (MOT) are two modern theories of chemical bonding. These are based on the wave mechanical model of atom.

􀁺 According to the valence bond theory the process of chemical bond formation can be visualised as the overlapping of atomic orbitals of the two atoms as they approach each other. The overlap increases the electron charge density in the inter-nuclear region.

􀁺 In order to explain bonding in molecules containing more than two atoms, Pauling proposed the concept of hybridisation. In hybridisation, the atomic orbitals of the valence shell of the central atom ‘ hybridise’ or merge and give newer orbitals with proper orientations, which explain the shape of the molecule.

􀁺 According to the Molecular orbital theory the atomic orbitals of comparable energies and of suitable symmetry combine to give rise to an equal number of molecular orbitals. These molecular orbitals extend over the entire region of the molecule i.e., these are delocalised over the whole molecule.

􀁺 When two atomic orbitals combine it gives a pair of molecular orbitals; one is called bonding molecular orbital of lower energy and the other of higher energy is called anti-bonding orbital.

􀁺 The electrons present in the molecule are filled in these orbitals in the order of increasing energy (Aufbau principle) to give the MO electronic configuration.

􀁺 The number of bonds between the two atoms is called bond order and is defined as Bond order = b.o. = ½(nb - na)

􀁺 The MO electronic configuration can be used to predict the magnetic nature of the molecule. If all the MO’s are doubly occupied the substance shows diamagnetic behaviour and if one or more MO’s are singly occupied the substance shows paramagnetic behaviour.

TERMINAL EXERCISE

1. What do you understand by a chemical bond?

2. Explain the process of bond formation as a decrease in energy.

3. What do you understand by the term, ‘bond length’ ?

4. Describe the two possible ways in which the noble gas electronic configuration is achieved in the process of bond formation.

5. What are Lewis electron dot symbols ? Show the formation of MgCl2 in terms of Lewis symbols.

6. Define a coordinate bond and give some examples.

7. What is VSEPR theory ? predict the shape of SF6 molecule using this theory.

8. Why do we need the concept of hybridisation ? How does it help in explaining the shape of methane ?

9. Give the salient features of molecular orbital theory.

10. Be2 molecule does not exist. Explain on the basis of molecular orbital theory.

11. Write down the molecular orbital electronic configuration of the following species and compute their bond orders.

O2 ; O2

+ ; O2

− ; O2

2−

12. BF3 is a polar molecule but it does not show dipole moment. Why?

13. Atom A and B combine to form AB molecule. If the difference in the electronegativity between A and B is 1.7. What type of bond do you expect in AB molecule?

14. Write down the resonating structures of N2O, SO4

2–, CO3

2– and BF3.