This is special case of dipole-dipole interaction.. This is found in the molecules in which highly polar N-H, O-H or H-F bonds are present. Although hydrogen bonding is regarded as being limited to N, O and F; but species such as Chlorine may also participate in hydrogen bonding. Energy of hydrogen bond varies between 10 to 100 kJ mol.

Hydrogen bond can be classified in three types

1. Inter molecular H bond. Same molecule

  1. Inter molecular H bond. Different molecule

Hydrogen Bonding between Ammonia and Water

  1. Intra Molecular Hyrogen Bond


Hydrogen bonds can vary in strength from weak (1–2 kJ mol−1) to strong (161.5 kJ mol−1 in the ion Typical enthalpies in vapour include:

F − H···: F (161.5 kJ/mol or 38.6 kcal/mol), illustrated uniquely by HF2,

O − H···: N (29 kJ/mol or 6.9 kcal/mol), illustrated water-ammonia

O − H···: O (21 kJ/mol or 5.0 kcal/mol), illustrated water-water, and alcohol-alcohol

N − H···: N (13 kJ/mol or 3.1 kcal/mol), illustrated by ammonia-ammonia

N − H···: O (8 kJ/mol or 1.9 kcal/mol), illustrated water-amide

Bond Length

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 always determined experimentally

A. Ionic Compound:- bond length is sum of cationic radius and anionic radius.

B. Covalent compound:

The bond length is measured by spectroscopic, x-ray diffraction and electron diffraction. Covalent radius is then calculated from this.

Lewis structure:- unable to give bond angle

VESPR they:- much better in predicate bond angle

M. O. T:- Give bond angle much better

Bond Energy or 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.

For a diatomic molecular Bond energy simply energy required to break the bond.

For Hetero-nuclear molecular it is average energy or mean energy.

Ex.   H2O(g) → H(g) + OH(g)                  ∆H = 502

OH(g) → H(g) + O(g)                              ∆H = 427

Bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types, relatively independent of the rest of the molecule


1. Sigma bond (s )> Pie (p) bond

2. Triple bond > double bond > single bond

3. s – s overlap > s – p overlap > p – p overlap


Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding. According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an eight electrons in their valence shells.

Exceptions to the Octet Rule:

  1. Hydrogen molecule: Hydrogen has one electron in its first energy shell (n = 1). It needs only one more electron to fill this shell, because the first shell cannot have more than two electrons. This configuration (1s2) is similar to that of noble gas helium and is stable. In this case, therefore, octet is not needed to achieve a stable configuration

Incomplete octet of the central atom: The octet rule cannot explain the formation of certain molecules of lithium, beryllium, boron, aluminium, etc.  (LiCl, BeH2, BeCl2, BH3, BF3) in which the central atom has less than eight electrons in the valence shell as shown below:

Expanded octet of the central atom: There are many stable molecules which have more than eight electrons in their valence shells. For example, PF5, has ten; SF6 has twelve and IF7 ha fourteen electrons around the central atoms, P, S, and I respectively  

Odd electron molecules: There are certain molecules which have odd number of electrons, like nitric oxide, NO and Nitrogen dioxide, NO2. In these cases, octet rule is not satisfied for all the atoms.

 It may be noted that the octet rule is based upon the chemical inertness of noble gases. However, it has been found that some noble gases (especially xenon and krypton) also combine with oxygen and fluorine to form a large number of compounds such a XeF2, KrF2, XeOF2, XeOF4, XeF6, etc.

This theory does not account for the shape of the molecules. It cannot explain the relative stability of the molecule in terms of the energy.


    Lewis structure and bonding theory cannot explain the shape of molecule.

Shape of H2O.

H → 1s2

O → 1s2 2s2 px2 py1 pz1

Angle H ─ O ─ H

Is 90o as per this mode but shell L = 104.5o


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