Chemical Bonding and Molecular Structure Class 11 Notes Chemistry Chapter 4

Introduction, Chemical Bond, Lewis Dot Structures, Ionic Bond, Covalent Bond, Valence Bond Theory (VBT), Hybridisation, Resonance, Hydrogen Bonding


In the previous chapter we have discussed about The Classification of Elements and Periodicity in Properties but in this chapter we will study Chemical Bonding and Molecular Structure.

Structure and Bonding is the heart of chemistry. Chemical bond is very important to explain the properties and structure of compound. The important aspect of each type of force is its relative strength, how rapidly it decreases with increasing distance and whether it is directional in nature or not.

Chemical Bond

It is the force of attraction between two atoms which hold them together in a compound or molecule. Nature loves stability and bond formation is associated with stability. Every element has a tendency to occupy inert electronic configuration which is considered as very stable. Noble gas electronic configuration can be achieved by

  1. Transference of electrons
  2. Mutual sharing of electrons
  3. Donation of lone pair of electrons

Types of Bond

In order to explain the formation of a chemical bond in terms of electrons, Lewis postulated that atoms achieve stable octet when they are linked by a chemical bond. On the basis of this chemical bonds are following type :

  1. Ionic bond
  2. Covalent bond
  3. Co-ordinate bond
  4. Metallic bond
  5. Hydrogen bond
  6. van der Waal’s bond

Lewis Dot Structures

Valence Electrons

In the formation of a molecule only the outer shell electrons take part in chemical bond combination and they are known as valence electrons. In Lewis symbols, an element is shown with symbol and valence electrons.

Octet Rule

It is proposed by Kossel and Lewis and according to this, "Every atom has a tendency to attain Noble gas electronic configuration or to have 8 valence electrons". This is known as law of octet rule or if it has two valence electrons then this is known as law of duplet. According to Lewis, only those compounds will be stable which follow octet rule.

Read also: States of Matter Class 11 Notes Chemistry Chapter 5

Formal Charge

Formal charge on an atom is the difference between the number of valence electrons is an isolated atom and the number of electrons assigned to that atom in a Lewis structure. It is expressed as:

Chemical Bonding and Molecular Structure Class 11 Notes

Ionic Bond

An ionic bond is formed by complete transference of one or more electrons from the valence shell of one atom to the valence shell of another atom. In this way both the atoms acquire stable electronic configurations of noble gases. The atom which loses electron becomes a positive ion and the atom which gains electron becomes negative ion.

Chemical Bonding and Molecular Structure Class 11 Notes

Note : Electrovalency is the number of electrons lost or gained during the formation of an ionic bond or electrovalent bond.

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Characteristics of Ionic Compounds

  1. They are hard, brittle and crystalline.

  2. They have high melting and boiling points.

  3. They are polar in nature.

  4. The linkage between oppositely charged ions is non rigid and non directional.

  5. They are soluble in polar solvents such as water and insoluble in non polar solvents such as CCl4, Benzene, ether etc.

  6. They are good conductors of electricity in fused state and in solution due to mobility of the ions. They are bad conductors of electricity in solid state because ions are unable to move.

Covalent Bond

A force which binds atoms of same or different elements by mutual sharing of electrons is called a covalent bond. If the combining atoms are same the covalent molecule is known as homoatomic. If they are different, they are known as heteroatomic molecule.

Chemical Bonding and Molecular Structure Class 11 Notes

Valence Bond Theory (VBT)

Valence bond theory was introduced by Heitler and London (1927) and developed by Pauling and others. It is based on the concept of atomic orbitals and the electronic configuration of the atoms. Let two hydrogen atoms A and B having their nuclei NA and NB and electrons present in them are eA and eB. As these two atoms come closer new attractive and repulsive forces begin to operate.

  1. The nucleus of one atom is attracted towards its own electron and the electron of the other and vice versa.

  2. Repulsive forces arise between the electrons of two atoms and nuclei of two atoms. Attractive forces tend to bring the two atoms closer whereas repulsive forces tend to push them apart.

Chemical Bonding and Molecular Structure Class 11 Notes

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Orbital overlap concept

If we refer to the minimum energy state in the formation of hydrogen molecule the two H-atoms are enough near so as to allow their atomic orbitals to undergo partial interpenetration. This partial interpenetration of atomic orbitals is called overlapping of atomic orbitals. The overlap between the atomic orbitals can be positive, negative or zero depending upon the characteristics of the orbitals participating to overlap.

Chemical Bonding and Molecular Structure Class 11 Notes

Types of overlapping

The covalent bonds can be classified into two different categories depending upon the type of overlapping. These are:

1. Sigma (σ) bond

This type of covalent bond is formed by the axial overlapping of half-filled atomic orbitals. The atomic orbitals overlap along the internuclear axis and involve end to end or head on overlap. There can be three type of axial overlap among s and p-orbitals as discussed below:

(i) s-s overlap : In this case, there is overlap of two half-filled s-orbitals along the internuclear axis as shown below.

Chemical Bonding and Molecular Structure Class 11 Notes

(ii) s-p overlapping : It involves the overlapping of half filled s-orbitals of one atom with the half filled p-orbitals of the other atom. The bond thus formed is called s-p sigma bond.

Chemical Bonding and Molecular Structure Class 11 Notes

(iii) p-p overlapping : It involves the co-axial overlapping between half filled p-orbitals of one atom with half filled p-orbitals of the other atom. The bond as formed is called p-p sigma bond.

Chemical Bonding and Molecular Structure Class 11 Notes

2. pi (π) bond

This type of covalent bond is formed when the atomic orbitals overlap in such a way that their axis remain parallel to each other and perpendicular to the internuclear axis. The orbitals formed due to sidewise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms. The electrons involved in the π bond formation are called pi-electrons.

Chemical Bonding and Molecular Structure Class 11 Notes


Hybridisation is the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape. The atomic orbitals combine to form new set of equivalent orbitals known as hybrid orbitals.

Salient Features of Hybridisation

  1. The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.

  2. The hybridised orbitals are always equivalent in energy and shape.

  3. The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.

  4. The type of hybridisation indicates the geometry of the molecules.

Important conditions for hybridisation

  1. The orbitals present in the valence shell of the atom are hybridised.

  2. The orbitals taking part in hybridisation must have only a small difference of energies.

  3. Promotion of electron is not essential condition prior to hybridisation.

  4. It is not necessary that only half filled orbitals participate in hybridisation.

Types of hybridisation

There are many different types of hybridisation depending upon the type of orbitals involved in mixing such as sp3, sp2, sp, sp3d, sp3d2 etc.

(i) sp-hybridisation

In this hybridisation one s and one p orbitals hybridise to produce two equivalent hybrid orbitals, known as sp hybrid orbitals. The two sp-hybrid orbitals are oriented in a straight line making an angle of 180° and therefore the molecule possesses linear geometry. Each of hybrid orbitals has 50% s-character and 50% p-character.

Example of molecules having sp-hybridisation are BeF2, BeCl2, BeHetc.

Chemical Bonding and Molecular Structure Class 11 Notes

(ii) sp2-hybridisation

In this hybridisation one s and one 2p orbitals hybridise to produce three equivalent hybrid orbitals, known as sp2 hybrid orbitals. sp2 hybrid orbitals are larger in size than sp-hybrid orbitals but slightly smaller than that of sp3 hybrid orbitals. Each sp2 hybrid orbitals has 1/3 (or 33.33%) s-character and 2/3 (or 66.7%) p-character. Example, BF3, BCl3, BH3 etc.

Chemical Bonding and Molecular Structure Class 11 Notes

(iii) sp3-hybridisation

In this hybridisation one s and three p-orbitals intermix to form sp3 hybrid orbitals of equivalent energy and identical shape. These four sp3 hybrid orbitals are directed towards the four corners of a tetrahedron separated by an angle of 109° 28'. sp3 hybrid orbitals have 1/4 (or 25%) s-character and 3/4 (or 75%) p-character. Example, CH4, NH3etc.

Chemical Bonding and Molecular Structure Class 11 Notes

(iv) sp3d-hybridisation

This type of hybridisation involves mixing of one s, three p and one d-orbitals to form five sp3d hybridised orbitals which adopt trigonal bipyramidal.

Formation of PCl5 : The ground state electronic configuration of phosphorus is 1s2 2s2 2p6 3s2 3p3. Under the conditions of bond formation the 3s-electrons get unpaired and one of the electron is promoted to vacant 3dz2 orbital. The ground state and excited state configurations of phosphorus are shown below :

Chemical Bonding and Molecular Structure Class 11 Notes
Chemical Bonding and Molecular Structure Class 11 Notes

(v) sp3d2-hybridisation

In this type of hybridisation one s, three p and two d-orbitals undergo intermixing to form six identical sp3d2 hybrid orbitals. These six orbitals are directed towards the corners of an octahedron and lie in space at an angle of 90° to one another.

Formation of SF6 : The ground state outer configuration of 16S is 3s2 3p4. In the excited state the electron pairs in 3s and 3px orbitals get unpaired and one out of each pair is promoted to vacant 3dz2 and 3dx2-y2 orbitals. The ground state and excited state configuration of 16S are given as follows:

Chemical Bonding and Molecular Structure Class 11 Notes
Chemical Bonding and Molecular Structure Class 11 Notes

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Sidgwick and Powell in 1940, proposed a simple theory based on repulsive character of electron pairs in the valence shell of the atoms. It was further developed by Nyholm and Gillespie (1957). Main Postulates are the following:

  1. The exact shape of molecule depends upon the number of electron pairs (bonded or non bonded) around the central atoms.

  2. The electron pairs have a tendency to repel each other since they exist around the central atom and the electron clouds are negatively charged.

  3. Electron pairs try to take such position which can minimize the rupulsion between them.

  4. The valence shell is taken as a sphere with the electron pairs placed at maximum distance.

  5. A multiple bond is treated as if it is a single electron pair and the electron pairs which constitute the bond as single pairs.

Bond Parameters

(i). Bond Angle

It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. Bond angle is expressed in degree which can be experimentally determined by spectroscopic methods.

Chemical Bonding and Molecular Structure Class 11 Notes

(ii). Bond Length

Bond length is defined as the equilibrium distance between the nuclei of two bonded atoms in a molecule.

Chemical Bonding and Molecular Structure Class 11 Notes

(iii). Lattice Enthalpy

The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions. For example, the lattice enthalpy of NaCl is 788 kJ mol–1.

(iv). Bond Order

Bond order is defined as half of the difference between the number of electrons present in bonding and antibonding molecular orbitals. The bond order may be a whole number, a fraction or even zero. It may also be positive or negative.

Bond order (B.O.) = `1/2[N_{b}-N_{a}]`

(vi). Bond Enthalpy

It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state. The unit of bond enthalpy is kJ mol–1. For example, the H – H bond enthalpy in hydrogen molecule is 435.8 kJ mol–1.

Molecular Orbital Theory (MOT)

Molecular orbital (MO) theory was developed by F. Hund and R.S. Mulliken in 1932. According to MOT, a molecule is considered to be quite different from the constituent atoms. All the electrons belonging to the atoms constituting a molecule are considered to be moving along the entire molecule under the influence of all the nuclei. Thus, a molecule is supposed to have orbitals of varying energy levels, in same way as an atom. These orbitals are called molecular orbitals.

Energy Level Diagram for Molecular Orbitals

Chemical Bonding and Molecular Structure Class 11 Notes


When a compound has same molecular formula but different structural formulas and structures differ with respect to electrons only. These structures are known as resonating structures or canonical structures. None of these structures can explain all the properties of that compound. This phenomenon is known as resonance.

Chemical Bonding and Molecular Structure Class 11 Notes

Hydrogen Bonding

When highly electronegative elements like nitrogen, oxygen, flourine are attached to hydrogen to form covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom. Thus, partial positive charge develops on hydrogen atom which forms a bond with the other electronegative atom. This bond is known as hydrogen bond and it is weaker than the covalent bond.

Types of Hydrogen Bonding

There are two types of hydrogen bonding.

1. Intermolecular hydrogen bonding : It is a type of hydrogen bonding between two similar or dissimilar molecules. Example : H – F, HF and water, NH3, NH3 and water, alcohol, alcohol and water etc.

Chemical Bonding and Molecular Structure Class 11 Notes

2. Intramolecular hydrogen bonding : It is a type of hydrogen bonding within the molecule. Example : Salicylaldehyde, O-nitrophenol etc.

Chemical Bonding and Molecular Structure Class 11 Notes

Applications of Hydrogen Bonding

1. State : Hydrogen bonding may affect the state of a compound. For example, H2O is liquid at room temperature whereas H2S is gas. It is due to presence of intermolecular hydrogen bonding between H2O molecules, which is not present in H2S molecules.

2. Solubility : Only those covalent molecules are soluble in water which have tendency to form intermolecular hydrogen bonding with water molecules.

3. Boiling point : Intermolecular hydrogen bonding increases the boiling point of compound. For example, NH3 has higher boiling point than PH3. This is because, there is intermolecular hydrogen bonding in NH3 but not in PH3.

4. Density of ice is lower than water : In ice, hydrogen bonding gives rise to a cage like structure of H–O–H molecules, in which each H–O–H molecule is linked tetrahedrally to four other H–O–H molecule. In this structure, some vacant spaces are formed, which decrease the density of ice.

Metallic Bonding

The force that binds a metal atom to a number of electrons within its sphere of influence is known as metallic bond. This model could easily explain the following properties of metals:

  1. High electrical conductivity
  2. High thermal conductivity
  3. Bright metallic lustre
  4. Malleability
  5. Ductility
  6. Tensile strength
  7. Elasticity


  1. Chemical Bond : The force of attraction which holds various chemical entities in different species.

  2. Electrovalent Bond : The attractive force between the oppositely charged ions which comes into existence by the transference of electrons.

  3. Electrovalency : The number of electrons which an atom loses or gains while forming ionic or electrovalent bond.

  4. Covalent Bond : The bond comes into existence by the mutual sharing of electrons by the atoms participating in bonding.

  5. Valence Bond Approach of Covalent Bond : The bond is formed by the overlapping of halffilled atomic orbitals having electrons with opposite spins.

  6. Covalency : The number of half-filled atomic orbitals which an atom provides for participation in overlapping at the time of bonding.

  7. Dative Bond or Co-ordinate Bond : The bond is formed by sharing of electrons in which the shared pair of electrons is contributed by one of the atom called donor while the other atom is called acceptor.

  8. Hybridisation : The process of mixing or merging of orbitals (of slightly different energies) of an atom to form another set of orbitals with equivalent shape and energy.

  9. Geometry of the Molecule : The definite relative arrangement of the bonded atoms in a molecule.

  10. Regular and Irregular Geometry : The molecule is said to possess regular geometry if the repulsive interactions among the electron pair around the central atom are of equal magnitude. If the repulsive interactions among the electron pairs are unequal, the geometry is referred to as irregular.

  11. Electronegativity : The power of an atom to attract bonding pair of electrons towards itself.

  12. Dipole Moment (μ) : A vector quantity defined by the product of charge developed on any of the atom and distance between the atoms ; creating a dipole.

  13. Polar and Non-Polar Molecules : The molecules with dipole moment (μ) > 0 are called polar molecules while those with μ = 0 are non-polar molecules.

  14. Dipole-Dipole Interactions : The attractive interactions among the opposite ends of polar molecules in liquid and solid state.

  15. Hydrogen Bond : The electrostatic force of attraction between covalently bonded H-atom of one molecule and the electronegative atom (F or N or O) of the other molecule.

  16. Resonance : When a molecule is represented by more than one electronic arrangement none of which is able to explain the observed characteristics of the molecule, then the actual structure is intermediate of various electronic arrangements and is known as resonance hybrid. The various electronic arrangements are called resonating structures or canonical structure.

  17. Molecular Orbital Theory (MOT) : According to this theory, in molecules the electrons are present in new orbitals called molecular orbitals. Molecular orbitals are not associated with a particular atom but belong to nuclei of all the atoms constituting the molecule.

  18. LCAO Method : This is an approximate method, according to which the molecular orbitals are obtained by linear combination of atomic orbitals.

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