Chemical Bonding

Chemical bond:-  

When two atoms of same or different elements approach each other, the energy of the combination of the atoms becomes less than the sum of the energies of the two separate atoms at a large distance. We say that the two atoms have combined or a bond is formed between the two. The bond is called a chemical bond. Thus a chemical bond may be visualised as an effect that leads to the decrease in the energy. The combination of atoms leads to the formation of a molecule that has distinct properties different from that of the constituent atoms. 


Kossel and Lewis developed electronic theory of valence or theory of chemical bonding to explain the formation of chemical bond between the two atoms.

According to the electronic theory of valence, every atom tries to attain octet configuration (presence of eight electrons) in its valence shell by losing or gaining or by sharing of electrons. This is known as the "Octet Rule". The electrostatic forces of attraction that holds the two oppositely charged ions together are known as "electrovalent bond".

Atoms attain stability by the means of bond formation. The process of bond formation is associated with the lowering of energy of the system.Only valence electrons participate in chemical bonding, but not inner electrons.To understand the concept of valence electrons Lewis introduced the concept representing valence electrons with dots, which are called Lewis symbols. 

Lewis symbols:

Lewis Symbols: In the formation of a molecule, only the outer shell electrons take part in chemical combination and they are known as valence electrons. The inner shell electrons are well protected and are generally not involved in the combination process. G.N. Lewis, an American chemist introduced simple notations to represent valence electrons in an atom. These notations are called Lewis symbols.

Lewis symbol for Chlorine atom.

Kössel, in relation to chemical bonding, drew attention to the following facts:  

• In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by the noble gases;

• The formation of a negative ion from a halogen atom and a positive ion from an alkali metal atom is associated with the gain and loss of an electron by the respective atoms;

• The negative and positive ions thus formed attain stable noble gas electronic configurations. The noble gases (with the exception of helium which has a duplet of electrons) have a particularly stable outer shell configuration of eight (octet) electrons, ns2np6.

• The negative and positive ions are stabilized by electrostatic attraction.

For example, the formation of NaCl from sodium and chlorine, according to the above scheme, can be explained as:

The bond formed, as a result of the electrostatic attraction between the positive and negative ions was termed as the electrovalent bond. The electrovalence is thus equal to the number of unit charge(s) on the ion.

Octet Rule

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 octet in their valence shells. This is known as octet rule.

There are three types of chemical bonds: Ionic Bond, Covalent Bond, Co-ordinate Bond.

Ionic or Electrovalent Bond

Ionic or Electrovalent bond is formed by the complete transfer of electrons from one atom to another.. For Example, ,

(i) NaCl is an electrovalent compound. Formation of NaCl is given below:

Na+ ion has the configuration of Ne while Cl ion represents the configuration of Ar.

(ii) Formation of magnesium oxide from magnesium and oxygen.

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

Characteristic Properties of Ionic Compounds

  • These exist as crystalline solids in which the ions are arranged in a regular three dimensional structure. The ionic compounds are generally hard and brittle in nature.
  • These compounds have high melting and boiling points due to strong electrostatic interactions between the ions.
  • These are generally soluble in water and less soluble in non-polar solvents like ether, alcohol, etc.
  • These conduct electricity when in molten state or in aqueous solutions.

Factors Affecting the Formation of Ionic Bond

(i) Ionization enthalpy: As we know that ionization enthalpy of any element is the amount of energy required to remove an electron from outermost shell of an isolated gaseous atom to convert it into cation.
Hence, lesser the ionization enthalpy, easier will be the formation of a cation and have greater chance to form an ionic bond. Due to this reason alkali metals have more tendency to form an ionic bond.
Therefore, we can conclude that lower the ionization enthalpy, greater the chances of ionic bond formation.

(ii) Electron gain enthalpy (Electron affinities): It is defined as the energy released when an isolated gaseous atom takes up an electron to form anion. Greater the negative electron gain enthalpy, easier will be the formation of anion. Consequently, the probability of formation of ionic bond increases.
For example. Halogens possess high electron affinity. So, the formation of anion is very common in halogens.

(iii) Lattice energy or enthalpy: It is defined as the amount of energy required to separate 1 mole of ionic compound into separate oppositely charged ions.

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.

Lattice energy of an ionic compound depends upon following factors:

(i) Size of the ions: Smaller the size, greater will be the lattice energy.

(ii) Charge on the ions: Greater the magnitude of charge, greater the interionic attraction and hence higher the lattice energy.

Covalent Bond—

When the bond is formed between two or more atoms by mutual contribution and sharing of electrons, it is known as 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.
For Example,

Lewis Representation of Simple Molecules (the Lewis Structures)

  • Pair of bonded electrons is by means of a ‘dash’ (-) usually called a ‘bond’.
  • Lone pairs or ‘non-bonded’ electrons are represented by ‘dots’.
  • Electrons present in the last shell of atoms are called valence electrons.

Exceptions to the Octet Rule:

The octet rule, though useful, is not universal. It is quite useful for understanding the structures of most of the organic compounds and it applies mainly to the second period elements of the periodic table. There are three types of exceptions to the octet rule.

  • Species with odd number of electrons:

In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO2, the octet rule is not satisfied for all the atoms. 

  • Incomplete octet for the central atom: 

In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons. For examples-LiCl, BeH2 and BCl3

  • Expanded octet for the central atom:

Elements in and beyond the third period of the periodic table have, apart from 3s and 3p orbitals, 3d orbitals also available for bonding. In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet. Obviously the octet rule does not apply in such cases.

For examples-PF5, SF6 and H2SO

Formal Charge:

  • The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure. It is expressed as:

  • Formula Charge Calculation of SO42-: