CBSE Class 11 Chemistry Chemical Bonding And Molecular Structure Notes Set 07

These revision notes on Chemical Bonding and Molecular Structure for CBSE Class 11 Chemistry cover the octet rule, Lewis structures, formal charge, ionic bonding, bond parameters, VSEPR theory, hybridisation, molecular orbital theory and hydrogen bonding. The shape tables and orbital energy orders are given in a form that is easy to revise before tests.

Quick Revision Notes

1. Chemical Bond

It is defined as the force of attraction which holds various constituents (atoms, ions, etc.) together in different chemical species. The combining tendency of an element or ion is called its ‘valency’.
Factors due to which atoms interact with each other are
(i) Forces of attraction overpower the forces of repulsion due to decrease in energy.
(ii) Tendency to achieve nearest noble gas configuration.

2. Kossel-Lewis Approach to Chemical Bonding (Octet Rule)

According to this approach, the atoms of different elements take part in chemical combination in order to complete their octet (to have eight electrons in the outermost valence shell) or duplet (to have two valence electrons) in some cases such as H, Li, Be etc., or to attain the nearest noble gas configuration. This is known as octet rule.
The group valency of the elements is generally either equal to the number of dots in Lewis symbols or 8 minus the number of dots or valence electrons.

Limitations of octet rule are as follows
(i) In accordance to this rule, the shape of the molecule cannot be predicted.
(ii) The relative stability of molecule cannot be known by this rule.

Exceptions to Octet Rule
The octet rule is violated in a significant number of cases. These are :
(i) Electron deficient compounds : BeCl₂, BF₃, AlCl₃ etc.
(ii) Hypervalent compounds : PCl₅, SF₆, IF₇, H₂SO₄ etc.
(iii) Compounds of noble gases : XeF₂, XeF₆, XeF₄, KrF₂ etc.
(iv) Odd electron molecules : NO, NO₂, O₂⁻ etc.
(v) Others like : H₂⁺, He₂⁺, O₂, NO, NO₂, ClO₂ are some of the examples of stable molecules having odd electron bonds (bonds formed by sharing of usually one or three electrons).

3. Types of Chemical Bonds

The three types of bonds or linkages which hold the atoms together in a molecule are as follows
(i) Covalent bond It is formed by mutual sharing of electrons between two atoms so as to attain nearest noble gas configuration.
(ii) Electrovalent bond It is formed as a result of the electrostatic attraction between the positive and negative ions.
(iii) Coordinate bond It is a type of covalent bond in which shared pair of electrons come from one atom only.

4. Lewis Representation of Simple Molecules (Lewis Structures)

It is representation of valence electrons, done by dots around the symbol of the element. e.g. Nitrite ion (NO₂⁻)

5. Formal Charge on Atom or Ion

It is 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 structures.
Formal charge on an atom in a molecule/ion
= [total number of valence electrons in the free atom] − [total number of non-bonding (lone pair) electrons] − \( \frac{1}{2} \) [total number of bonding (shared) electrons]

6. Ionic or Electrovalent Bond

It is formed by complete transfer of electrons from one atom to another. e.g. Formation of sodium chloride

Factors governing the formation of ionic bond are
(i) Low ionisation enthalpy of metal atom.
(ii) High electron gain enthalpy of non-metal atom.
(iii) High lattice enthalpy of the ionic compound.

Lattice enthalpy
It is defined as the amount of energy required to completely separate one mole of a solid ionic compound into its gaseous constituent ions. It depends on
(i) charge on ions (greater is the charge, greater is the lattice enthalpy).
(ii) size of the ion (smaller is the size, greater is the lattice enthalpy).

7. Bond Parameters

(i) Bond length The equilibrium distance between the nuclei of two bonded atoms in a molecule is called the bond length.
The various factors that affect bond length are as follows
(a) Size of the atoms The bond length increases with increase in the size of the atoms. Thus, the order of bond length of H—X bond is HF < HCl < HBr < HI.
Similarly, C—C < Si—Si < Ge—Ge
(b) Bond multiplicity The bond length decreases as the multiplicity of the bond increases.
Thus, the carbon-carbon bond lengths follow the order C≡C < C=C < C—C

(ii) Bond angle The angle between the orbitals containing bonded electron pairs around central atom in a molecule/complex ion is called the bond angle.
e.g. H—C—H bond angle in CH₄ and H—N—H bond angle in NH₃ are shown below :

(iii) Bond enthalpy It is the amount of energy required to break one mole of bonds of a particular type between two atoms in gaseous state. Its unit is kJ mol⁻¹.
Bond enthalpy ∝ \( \frac{1}{\text{Size of atoms}} \) ∝ \( \frac{1}{\text{Number of lone pair(s) of electrons}} \) ∝ Multiplicity of bond
Greater the multiplicity of the bond, greater is the bond enthalpy.
e.g. H—H < O=O < N≡N

(iv) Bond order The number of bonds formed between the two atoms in a molecule is called the bond order.
e.g. The bond order of a few molecules are given below:

Bond order ∝ Bond enthalpy ∝ \( \frac{1}{\text{Bond length}} \)

(v) Resonance structures According to the concept of resonance, whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, positions of nuclei, bonding and non-bonding pairs of electrons are taken as the canonical structures of the hybrid which describes the molecule accurately.

(vi) Dipole moment It is a measurement of polarity in any molecule/compound/ion. It is defined as the product of charge (q) and distance between the charges (r).
Dipole moment (μ) = q × r. Its unit is Debye (D). The dipole moment of symmetrical molecule is always zero.

8. Partial Ionic Character of Covalent Bond

(i) Electronegativity difference between combining atoms = 1.7, then bond is 50% ionic and 50% covalent.
(ii) Electronegativity difference > 1.7, ionic character in bond is more than 50%.
(iii) Electronegativity difference < 1.7, ionic character in bond is less than 50%.

9. Partial Covalent Character in Ionic Bonds : Fajans’ Rule

According to Fajans’ rule, the covalent character will be favoured by
(i) small size of cation.
(ii) large size of anion.
(iii) high charge of cation and anion.
(iv) the cation having 18-electrons shell.

10. Valence Shell Electron Pair Repulsion (VSEPR) Theory

(i) According to this theory, all valence shell electron pairs surrounding the central atom arrange themselves in such a manner that they are as far away from each other as possible.
(ii) There are two types of electron pairs around the central atom; bonding electron pairs, i.e. bond pair (bp) and non-bonding electron pairs, i.e. lone pair (lp).
(iii) The strength of repulsion between the electron pairs varies as lp-lp > lp-bp > bp-bp
(iv) If central atom is linked to similar atoms and there are no lone pairs, the shape is symmetrical otherwise irregular.
(v) If A represents central atom, B bond pair and E lone pair, then AB₂ → linear.

Shape of Some Molecules or Ions with Central Atom having One or More Lone Pairs of e⁻ on the Basis of VSEPR Theory

11. Valence Bond Theory (VBT)

It explains bond formation in terms of overlapping of orbitals, e.g. the formation of H₂ molecule from two hydrogen atoms involves the overlapping of 1s-orbital of two H-atoms which are singly occupied. Because of orbital overlapping, the electron density between the nuclei increases which helps in bringing them closer. The overlapping of orbitals may result in two types of bonds, sigma (σ) and pi (π) bond.

12. Concept of Hybridisation (involving s, p, d and f-orbitals)

(i) Intermixing of orbitals of almost similar energy belonging to the same atom to form same number of orbitals of exactly equal energy is known as hybridisation.
(ii) The new orbitals thus formed are known as hybrid orbitals. The number of hybrid orbitals of the molecule can be determined by
\( H = \frac{1}{2}(V + Y - C + A) \)
where, V = valence shell electrons of the central atom
Y = number of monovalent atoms
C = total positive charge
A = total negative charge on the molecule

Geometry and Hybridisation of Molecules

13. Molecular Orbital Theory (MOT)

(i) Bonding molecular orbitals The molecular orbitals obtained by the addition of atomic orbitals are called bonding molecular orbitals (BMOs) and are represented by σ and π.
(ii) Anti-bonding molecular orbitals The molecular orbitals obtained by the subtraction of atomic orbitals are called anti-bonding molecular orbitals (AMOs) and are represented by σ* and π*.
The electronic configuration of molecular orbitals (MO) are written in the following manner.
(a) In the case of O₂ and F₂,
σ1s < σ*1s < σ2s < σ*2s < σ2p_z < π2p_x = π2p_y < π*2p_x = π*2p_y < σ*2p_z
(b) However, the above sequence of energy levels is not true in case of remaining molecules, like Li₂, Be₂, B₂, C₂, N₂. For these molecules, the increasing order of energies of various molecular orbitals is
σ1s < σ*1s < σ2s < σ*2s < π2p_x = π2p_y < σ2p_z < π*2p_x = π*2p_y < σ*2p_z

Stability and Magnetic Character of Molecules/Molecular Ions
(i) If bond order is positive, molecule or ion is stable otherwise not. Greater the bond order, greater is the stability and shorter is the bond length.
(ii) Further, if all electrons are paired, the species is diamagnetic and if some unpaired electrons are present, the species is paramagnetic.
(iii) Bond order = \( \frac{1}{2} \) [(Number of electrons in BMOs) − (Number of electrons in AMOs)]
i.e. BO = \( \frac{1}{2}(N_b - N_a) \)

14. Hydrogen Bonding

(i) The attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) of another molecule, is called hydrogen bonding.
(ii) It is represented by dotted lines, e.g. …H^δ+—F^δ−……H^δ+—F^δ−……H^δ+—F^δ−

Types of H-Bonds
There are two types of H-bonds :
(i) Intermolecular hydrogen bond Such bond exists between two different molecules of the same or different compounds.
(ii) Intramolecular hydrogen bond Such bond exists when hydrogen atom is internally bonded through H-bond with highly electronegative (F, O, N) atoms present within the same molecule.