Chemical Bonding and Molecular Structure

Kossel-Lewis Concept: Atoms take part in chemical combinations to complete octets in their valence shell. This is known as the octet rule.
Limitation of Octet Rule: The octet rule, though useful but have some exceptions e.g. BF₃, NO₂, PCl₅, SF_6, etc.

Lewis Symbol or Electron Dot Structure: Representing valence electrons by dots placed around the letter symbol of the element.

Types of Chemical Bonds:
(i) Covalent Bond:
(a) Formed by sharing of electrons.
(b) It may be polar or nonpolar.
(c) It is directional in nature.
(ii) Ionic Bond:
(a) Formed by transfer of electrons.
(b) Formation of an ionic bond is favored by high lattice enthalpy, Low ionization enthalpy of the metal atom, and more negative electron gain enthalpy of the nonmetal atom.
(c) It is non-directional in nature.

Formal Charge (F.C.):
(i) It is a charge that appeared on an individual atom in the covalent molecule.
(ii) F.C. = (Total No. of valence electrons in a free atom) – (Total No. of unshared electrons) – ½ (Total No. of shared electrons).
Greater the F.C on atoms lesser the stability of that Lewis structure.

Lattice Enthalpy: Energy released when one mole of a crystalline solid is formed constituent gaseous ions.
Bond length:
(i) It is the equilibrium distance between the nuclei of two bonded atoms in a molecule.
(ii) Greater the size of bonded atoms shorter the bond length. e.g., H–F < H–Cl < H–Br < H–I
(iii) Greater the s character shorter the bond length. e.g., C_sp3 –H > C_sp2 –H > C_sp –H >
(iv) Bond length decreases with an increase in bond order. e.g., C–C > C=C > C≡C.

Bond angle:
(i) It is the angle between the orbitals containing bonding electron pairs around the central atom in a molecule or complex ion.
(ii) Greater the electronegativity of the central atom larger the bond angle e.g., NH₃ > PH₃
(iii) Greater the number of lone pairs around the central atom smaller the bond angle. e.g., CH₄ > NH₃ > H₂O.

Bond Enthalpy:
(i) 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.
(ii) For diatomic molecules, Bond enthalpy = Bond dissociation enthalpy
(iii) For polyatomic molecules, Bond enthalpy = Average of all possible bond dissociation enthalpies.
(iv) Bond enthalpy α Bond order α 1/(Bond length).

Resonance:
(i) According to the concept of resonance, whenever a single Lewis structure cannot describe a molecule accurately, a number of structures with similar energy, the position of nuclei, bonding and non-bonding pairs of electrons are taken as canonical structures of the resonance hybrid which describes the molecule accurately.
(ii) Resonance averages the bond characteristics as a whole.

The partial ionic character of covalent bond A–B: = 16(X_A – X_B) + 3.5(X_A – X_B)², where X_A and X_B are electro-negativities of A & B.
Partial covalent character in ionic bond (Fajan’s rule):
(i) Fajan’s rule is used to predict partial covalent character in ionic bonds.
(ii) Greater the polarizing power of cation and polarisability of anion greater the covalent character in an ionic bond.
(iii) Polarising power of cation α Charge density [(Charge)/Radius)].
(iv) Polarisability of anion α size of the anion.

Chemical Bonding

Dipole moment:
(i) Dipole moment (μ) = charge (Q) × distance of separation (d)
(ii) Unit: Debye (D), 1D = 3.33564 × 10^–30 Cm
(iii) Being vector quantity, the dipole moment of the polyatomic molecule is taken as the resultant of all the bond moments.
(iv) If μ= 0, molecule is non polar or symmetric. (v) If μ≠ 0, the molecule is polar or asymmetric.

Hydrogen bond:
(i) It is a dipole-dipole interaction between molecules in which the ‘H’ atom is inserted between two highly electronegative elements i. e. F, O, or N only.
(ii) Hydrogen bonds may be intra-molecular (when present within a single molecule) and intermolecular (when present b/w two same or different molecules).
(iii) Hydrogen bonds are stronger intermolecular forces than van der Waal forces.

Sigma (σ) and pi (π) bonds:
(i) Sigma bond is formed by axial overlapping and the pi bond is formed by sideways overlapping of atomic orbitals.
(ii) Sigma bond is stronger than the pi bond due to the greater extent of overlapping.
(iii) Single covalent bond = 1 σ bond Double covalent bond = 1 σ bond + 1 π bond Triple covalent bond = 1 σ bond + 2 π bond.

VSEPR Theory: (VSEPR = Valence Shell Electron Pair Repulsion):
The shape of a molecule depends upon the number of valence shell electron pairs (lp and bp) around the central atom and the magnitude of repulsive forces between them i.e., lp–lp > lp–bp > bp–bp

Hybridization:
(i) It is the phenomenon of mixing of atomic orbitals of nearly the same energy to form new orbitals of equal energy and identical shape.
(ii) The new orbitals are called hybrid orbitals and determine the shape of molecules.

Molecular Orbital Theory (MOT):
(i) The overlap of atomic orbitals of the same symmetry to form bonding and antibonding molecular orbitals by addition and substructions of their wave functions is known as MO theory.
(ii) The electrons are filled in molecular orbitals in order of their increasing energy. i.e.,
σ1s, σ*1s, σ2s, σ*2s, π2px = π2py, σ2pz, π*2px = π*2py, σ*2pz (up to 14 electrons)
σ1s, σ*1s, σ2s, σ*2s, σ2pz, π2px = π2py, π*2px = π*2py, σ*2pz (For more than 14 electrons)
(iii) Bond order = 1/2 (N_b – N_a)
N_a = No of electrons in anti-bonding molecular orbitals
N_b = No of electrons in bonding molecular orbitals.