Explain ${\mathrm{sp}}^2$ hybridisation in ${\mathrm{BF}}_3$.

1. Kossel-lewis approach to chemical bonding:

(i) The force which holds the atoms together in a molecule is called a chemical bond.

(ii) Atoms combine so as to complete their octets ($8$ electrons in the outermost shell) or duplet ($2$ electrons in the outermost shell) in case of $\mathrm{H}, \mathrm{Li}, \mathrm{Be},$ etc. to attain nearest noble gas configuration and gain stability.

(iii) Atoms combine by:

(a) By transference of electrons from one atom to another (ionic bond/electrovalent bond)

(b) By mutual sharing of electrons (covalent bond)

(c) By sharing of electrons contributed by one atom (coordinate bond/dative bond)

2. Lewis symbol:

Lewis symbol is representation of valence electrons by dots around the symbol of the element.

3. Formal charge:

Formal charge on an atom in a molecule/ion $=$ Total no. of valence electrons in free atom $-$Total no. of electrons of lone pairs -12 (Total no. of shared electrons).

4. Valence Shell Electron Pair Repulsion (VSEPR) theory:

(i) It states, “Electron pairs surrounding the central atom repel one another and go far apart from one another until there are no further repulsions.”

(ii) The magnitude of repulsions is: Lone-Pair-Lone Pair $>$ Lone Pair-Bond Pair $>$ Bond Pair-Bond Pair.

(iii) If central atom is linked to similar atoms and there are no lone pairs, the shape is symmetrical (regular) otherwise irregular.

(iv) If A represents central atom, B bond pair and L lone pair, we have ${\mathrm{AB}}_2$(linear), ${\mathrm{AB}}_3$ (triangular planar), ${\mathrm{AB}}_4$ (tetrahedral), ${\mathrm{AB}}_2\mathrm{L}$ (V-shape), ${\mathrm{AB}}_2{\mathrm{L}}_2$ (bent), ${\mathrm{AB}}_3\mathrm{L}$ (trigonal pyramidal), ${\mathrm{AB}}_3{\mathrm{L}}_2$ (T-shape), ${\mathrm{AB}}_2{\mathrm{L}}_3$ (linear), ${\mathrm{AB}}_5$ (trigonal bipyramidal), ${\mathrm{AB}}_4\mathrm{L}$ (see-saw), ${\mathrm{AB}}_6$ (octahedral), ${\mathrm{AB}}_5\mathrm{L}$ (square pyramidal), and ${\mathrm{AB}}_4{\mathrm{L}}_2$ (square planar).

5. Calculation of total no. of bond pairs (B), electron pairs (T), and lone pairs (L):

$\mathrm{T}=\frac{1}{2}$ (Number of valence electrons of central atom $+$ Number of atoms linked to it by single bonds) $+$ No. of electrons equal to negative charge for anion or minus no. of electrons equal to positive charge for cation).

$\mathrm{B} =$ No. of atoms linked to central atom by single bonds

$\mathrm{L} = \mathrm{T}–\mathrm{B}$.

6. Sigma and pi bonds:

(i) When the overlap takes place along the internuclear axis, the bond formed is called sigma (σ) bond. It can be $\mathrm{s}–\mathrm{s}, \mathrm{s}–\mathrm{p}$ or $\mathrm{p}–\mathrm{p}$ overlap.

(ii) When a covalent bond is formed by sideways overlapping of orbitals, it is called pi $\left(\mathrm{\pi }\right)$ bond, e.g., between $\mathrm{px}$ and $\mathrm{px}$ or $\mathrm{py}$ and $\mathrm{py}$ orbitals.

(iii) $\mathrm{\sigma }$-bond is always stronger than $\mathrm{\pi }$ bond because overlapping is greater along the internuclear axis than sideways overlapping.

7. Polar and Non-polar bond:

(i) Covalent bond formed between two dissimilar atoms is called a polar bond.

(ii) Covalent bond formed between two similar atoms is called a non-polar bond.

8. Bond parameters:

Important parameters, associated with chemical bonds, like bond length, bond angle, bond enthalpy, bond order and bond polarity have significant effects on the properties of compounds.

9. Dipole moment:

It is the product of magnitude of charge (q) on each end and distance (d) between the centres of the nuclei of two bonded atoms, i.e.,

$\mathrm{\mu }=\mathrm{q}\mathrm{ }\times \mathrm{ }\mathrm{d}$. Its units are Debye.

10. Hybridisation:

Mixing of atomic orbitals of an atom with slightly different energies to form the same number of new orbitals of equal energy and identical shape is called hybridisation. The new orbitals formed are called hybrid orbitals.

11. Resonance:

When one Lewis structure of a molecule/ion cannot explain all of its properties, it is believed to have a number of structures, each of which can explain most of its properties but none of them can explain all of its properties. The actual structure is in between all these structures and is called resonance hybrid. The phenomenon is called resonance.

12. Molecular orbital theory (MOT):

(i) Molecular orbital theory (MOT) is an alternative approach to explain the nature and formation of covalent bonds. This theory uses all the electrons of the combining atoms to form molecular orbitals. Number of molecular orbitals formed is the same as the number of atomic orbitals. Two types of molecular orbitals are formed: bonding molecular orbitals and antibonding molecular orbitals.

(ii) The principle of Linear Combination of Atomic Orbitals (LCAO) is used to form the molecular orbitals.

(iii) Using molecular orbital theory, one can predict the magnetic character of the molecule or ion, bond order and variations in bond strengths and bond lengths.