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Chemical Bonding And Molecular Structure

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Chemical Bonding and Molecular Structure PDF Notes, Important Questions and Formulas

Chemical Bonding and Molecular Structure


A molecule will be formed if it is more stable and has a lower energy than the individual atom. Normally only electrons in the outermost shell of an atom are involved in forming bonds, and in this process each atom attains a stable electronic configuration of an inert gas.

Types of Bonds:

Atoms Involved


I.   Metal + Nonmetal


II.  Nonmetal + Nonmetal


III. Metal + Metal


IV. Electron deficient molecule or ion (Lewis acid) and electron rich molecule or ion           (Lewis base)


V.  H and electronegative elements (N, O, F sometimes with C1)


Ionic, covalent and coordinate bonding Ionic bonding

An ionic bond is formed when a metal atom transfers one or more electron to a non-metal atom.

The oppositely charged ions get attracted through electrostatic force of attraction.

Properties of ionic compounds:

  • In the solid state, each cation surrounds itself with anions and each anion with cations. These very large number of ions are arranged in an ordinary network called ionic crystals.
  • They are good conductors of electricity in fused state or aqueous solution.
  • They are generally soluble in polar solvents and insoluble in non-polar solvents.
  • Have high melting point and boiling point than covalent compounds.
  • Have strong force of attraction between cation and anion (Coulombic force)

Covalent Bonding:

Whenever Chemical bond is formed by sharing of electrons then it is named as covalent bond.

Coordinate Bonding or Dative Bond:

Whenever covalent bond is formed by unequal sharing of a pair of electrons between a Lewis base and Lewis acid is called coordinate bond.

  1. It is represented as (→) and considered as 𝜎 bond.
  2. Atom/ion/molecule donating electron pair is called donor or Lewis base.
  3. Atom/ion/molecule accepting electron pair is called acceptor or Lewis acid.

Valence Bond (V.B) theory:

According to this theory, a covalent bond is formed by the overlapping of atomic orbitals. Important points of this theory are summarised below.

  1. Orbitals undergoing overlapping should be half-filled.
  2. Half-filled orbitals should contain the electron with opposite spin.
  3. Strength of a covalent bond depends upon the extent of overlapping, for example, axial or lateral overlapping.
  4. If the atomic orbitals overlap axially, then the bond formed is called a sigma (𝜎 ) bond,
  5. A sigma bond is always stronger than a pi-bond.
  6. Covalent bonds formed by the overlap of s-s and s-p orbitals are always sigma.
  7. By the overlap of p-p orbitals, one sigma and two pi bonds are formed.
  8. Increasing strength of 𝜎 covalent bonds is in the order s – s< s – p < p –p (when internuclear distance is constant )
  9. A single covalent bond is always a sigma bond. In a double covalent bond, one is sigma and the other is pi-bond. In a triple covalent bond, one is sigma and two are pi-bonds.


It is defined as the concept of intermixing of orbitals of equivalent energy, identical shapes and which are symmetrically disposed in plane. Important features of hybridisation are given below.

  1. Only the orbitals generated are equal in number to that of pure atomic orbitals which are intermixed.
  2. The hybrid orbitals generated are equal in number to that of pure atomic orbitals which are intermixed
  3. A hybrid orbitals, like an atomic oprbital, can have two electrons of opposite spins.
  4. Hybrid orbitals usually form sigma bonds. If there are pi-bonds, equal    number of atomic orbitals must be left unhybridised for pi-bonding.

Valence shell electron pair repulsion (VSEPR) Theory (Gillespie theory):

The shape of a molecule is determined by repulsion between the electron pairs begin mathsize 12px style left parenthesis calligraphic l straight p text  and bp end text right parenthesis end stylepresent in the valence shell of the central atom.

The order of repulsion isbegin mathsize 12px style left parenthesis calligraphic l straight p minus calligraphic l straight p right parenthesis greater than left parenthesis calligraphic l straight p minus bp right parenthesis greater than left parenthesis bp minus bp right parenthesis end style 

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