c) assertion is true but reason is false

In the case of covalent bond formed between atoms having different electronegativities, the atom with higher electro negativity will have greater tendency to attract the shared pair of electrons more towards itself than the other atom. As a result the cloud of shared electron
pair gets distorted. Let us consider the covalent bond between hydrogen and fluorine in hydrogen fluoride. The electro negativities ofhydrogen and fluorine on Pauling's scale are 2.1 and 4 respectively. It means that fluorine attracts the shared pair of electrons approximately twice as much as the hydrogen which leads to partial negative charge on fluorine and partial positive charge on hydragen. Hence, the H-F bond is said to be polar covalent bond.

H₂ molecule :
Electronic configuration of H atom 1s¹
Electronic configuration of H₂ molecule \(\sigma^{2}_{1s}\)
Bond order = \({N_b-N_a\over2}={2-0\over2}=1\)
Molecule has no unpaired electrons hence it is diamagnetic

Helium molecule (i.e) He₂ :
The electronic econfiguration of He atom is 1 s²
.'. Electronic configuration of He₂ molecule is \((\sigma_{1s})^2(\sigma^*_{1s})^2\)
Bond order = \({N_b-N_a\over2}={2-2\over2}\)
= 0
He, cannot exist. Similarly all the inert gases, X₂ cannot orist as their bond order = 0.

Bonding in some heteronuclear di-atomic molecules:
Molecular orbital diagram of Carbon monoxide molecule (CO)
Electronic configuration of C atom 1s²2s²2p²
Electronic configuration of O atom 1s²2s²2p⁴
Electronic conguration of CO molecule
\(\sigma^{2}_{1s},\sigma^{*2}_{1s},\sigma^{2}_{2s},\sigma^{*2}_{2s}\)\(\pi^{2}_{2py},\pi^{2}_{2pz},\sigma^2_{2px}\)
Bond order = \({N_b-N_a\over2}={10-4\over2}=3\)
Molecule has no unpaired electrons hence it is diamagnetic

Molecular orbital diagram of oxyge molecule (O₂) :
Electronic configuration of O atom 1s²2s²2p⁴
Electronic configuration of O₂ molecule
\(\sigma^{2}_{1s},\sigma^{*2}_{1s},\sigma^{2}_{2s},\sigma^{*2}_{2s},\sigma^{2}_{2px},\)\(\pi^{2}_{2py},\pi^{2}_{2pz},\pi^{*1}_{2py},\pi^{*1}_{2pz}\)
Bond order=\({N_b-N_a\over2}={10-6\over2}=2\)
Molecule has two unpaired electrons hence it is paramagnetic.

The ability of a cation to polarise an anion is called its polarising ability and the tendency of an anion to get polarised is called its polarisability. The extent of polarisation in an ionic compound is given by the Fajans rules
Fajans Rules:
(i) To show greater covalent character, both the cation and anion should have high charge on them. Higher the positive charge on the cation, greater will be the attraction on the electron cloud of the anion. Similarly higher the magnitude of negative charge on the anion, greater is its polarisability. Hence, the increase in charge on cation or in anion increases the covalent character Let us consider three ionic compounds aluminum chloride, magnesium chloride and sodium chloride. Since the charge of the cation increase in the order Na⁺ < Mg²⁺ < AI³⁺, the covalent character also follows the same order NaCI < MgCl₂ < AICI₃.
2. The smaller cation and larger anion show greater covalent character due to the greater extent of polarisation.
Lithium chloride is more covalent than sodium chloride. e size of Li⁺ is smaller than Na⁺ and hence the polarising power of Li⁺ is more. Lithium iodide is more covalent than lithium chloride as the size of I^- is larger than the CI^-. Hence I^- will be more polarised than CI^- by the cation, Li⁺.
3. Cations having ns² np⁶ nd¹⁰ configuration show greater polarising power than the cations with ns² np⁶ configuration. Hence, they show greater covalent character.
CuCI is more covalent than NaCl. Compared to Na+ (1.13 \(\mathring { A } \)). Cu+ (0.6 \(\mathring { A } \)) is small and have 3s² 3p⁶ 3d¹⁰ conguration.
Electronic conguration of Cu⁺
[Ar] 3s², 3p⁶, 3d¹⁰
Electronic Conguration of Na⁺
[He] 2s², p⁶

Molecules having polar bonds will not necessarily have a dipole moment. For example, the linear form of carbon dioxide haszero dipole moment, even though it has two polar bonds. In CO₂, the dipole moments of two polar bonds (CO) are equal in magnitude but have opposite direction. Hence, the net dipole moment of the CO₂ is,
\(\mu=\mu_{1}+\mu_{2}=\mu_{1}+(-\mu_{1})=0.\)

In this case \(\mu=\overrightarrow { { \mu }_{ 1 } } +\overrightarrow { { \mu }_{ 2 } } \quad \)
\(=\overrightarrow { { \mu }_{ 1 } } +\overrightarrow { ({ \mu }_{ 1 }) } =0\)
Incase of water net dipole moment is the vector sum of \(\mu_{1}+\mu_{2}\) as shown.

Dipole moment in water is found to be 1.85 D

Important principles of VSEPR Theory are as follows:
1. The shape of the molecules depends on the number of valence shell electron pair around the central atom.
2. There are two types of electron pairs namely bond pairs and lone pairs. The bond pair of electrons are those shared between two atoms, while the lone pairs are the valence electron pairs that are not involved in bonding.
3. Each pair of valence electrons around the central atom repels each other and hence, they are located as far away as possible in three dimensional space to minimize the repulsion between them.
4. The repulsive interaction between the different types of electron pairs is in the following order.
lp - lp > lp - bp > bp - bp
lp - lone pair; bp - bond pair
The lone pair of electrons are localised only on the central atom and interacts with only one nucleus whereas the bond pairs are shared between two'atoms and they interact with two nuclei. Because of this the ione pairs occupy more space and have greater repulsive Power than the bond Pairs in a molecule.

Molecular orbital diagram of nitrogen molecule (N₂) :
Electronic configuration of N atom 1s² 2s² 2p³
Electronic configuration of N, molecule
\(\sigma^{2}_{1s},\sigma^{*2}_{1s},\sigma^{2}_{2s},\sigma^{*2}_{2s}\)\(\pi^{2}_{2py},\pi^{2}_{2pz},\sigma^2_{2px}\)
Bond order = \({N_b-N_a\over2}={10-4\over2}=3\)
Molecule has no unpaired electrons hence it is diamagnetic.