First member of each group of representative elements (i.e., $s$ - and p-block elements) shows anomalous behaviour due to (i) small size (ii) high ionisation enthalpy (iii) high electronegativity and (iv) absence of $d$ - orbitals. e.g., in s-block elements, lithium shows anomalous behaviour from rest of the alkali metals.

(a) Compounds of lithium have significant covalent character. While compounds of other alkali metals are predominantly ionic.

(b) Lithium reacts with nitrogen to form lithium nitride while other alkali metals do not form nitrides.

In p-block elements, first member of each group has four orbitals, one $2 s$ - orbital and three $2 p$-orbitals in their valence shell. So, these elements show a maximum covalency of four while other members of the same group or different group show a maximum covalency beyond four due to availability of vacant $d$ - orbitals.

In p-block, when we move from left to right in a period, the acidic character of the oxides increases due to increase in electronegativity. e.g.,

(i) 2nd period $\mathrm{B}_2 \mathrm{O}_3<\mathrm{CO}_2<\mathrm{N}_2 \mathrm{O}_3$ acidic nature increases.

(ii) 3rd period $\mathrm{Al}_2 \mathrm{O}_3<\mathrm{SiO}_2<\mathrm{P}_4 \mathrm{O}_{10}<\mathrm{SO}_3<\mathrm{Cl}_2 \mathrm{O}_7$ acidic character increases.

On moving down the group, acidic character decreases and basic character increases. e.g.,

(a) Nature of oxides of 13 group elements

$$ \underset{\text { Weakly acidic }}{\mathrm{B}_2 \mathrm{O}_3} \underbrace{\mathrm{Al}_2 \mathrm{O}_3 \quad \mathrm{Ga}_2 \mathrm{O}_3}_{\text {Amphoteric }} \underset{\text { Basic }}{\mathrm{In}_2 \mathrm{O}_3} \underset{\text { Strongly basic}}{\mathrm{Tl}_2 \mathrm{O}} $$

(b) Nature of oxides of 15 group elements

$\underset{\text { Strongly acidic}}{\mathrm{N}_2 \mathrm{O}_5} ~\underset{\text { Moderately acidic}}{\mathrm{P}_4 \mathrm{O}_{10}} ~\underset{\text { Amphoteric}}{\mathrm{As}_4 \mathrm{O}_{10}}~ \underset{\text { Amphoteric}}{\mathrm{Sb}_4 \mathrm{O}_{10}} ~\underset{\text { Basic}}{\mathrm{Bi}_2 \mathrm{O}_3}$

Among the oxides of same element, higher the oxidation state of the element, stronger is the acid. e.g., $\mathrm{SO}_3$ is a stronger acid than $\mathrm{SO}_2$.

$\mathrm{B}_2 \mathrm{O}_3$ is weakly acidic and on dissolution in water, it forms orthoboric acid. Orthoboric acid does not act as a protonic acid (it does not ionise) but acts as a weak Lewis acid.

$$\underset{\text { Boron trioxide }}{\mathrm{B}_2 \mathrm{O}_3}+3 \mathrm{H}_2 \mathrm{O} \rightleftharpoons \underset{\text { Orthoboric acid }}{2 \mathrm{H}_3 \mathrm{BO}_3}$$

$$\mathrm{B}(\mathrm{OH})_3+\mathrm{H}-\mathrm{OH} \longrightarrow\left[\mathrm{B}(\mathrm{OH})_4\right]^{-}+\mathrm{H}^{+}$$

$\mathrm{Al}_2 \mathrm{O}_3$ is amphoteric in nature. It is insoluble in water but dissolves in alkalies and reacts with acids.

$$\underset{\begin{array}{c} \text { Aluminium } \\ \text { trioxide } \end{array}}{\mathrm{Al}_2 \mathrm{O}_3}+2 \mathrm{NaOH} \xrightarrow{\Delta} \underset{\begin{array}{c} \text { Sodium meta } \\ \text { aluminate } \end{array}}{2 \mathrm{NaAlO}_2}+\mathrm{H}_2 \mathrm{O} \longleftarrow \mathrm{Al}_2 \mathrm{O}_3+6 \mathrm{HCl} \xrightarrow{\Delta} \underset{\begin{array}{c} \text { Aluminium } \\ \text { chloride } \end{array}}{2 \mathrm{AlCl}_3}+3 \mathrm{H}_2 \mathrm{O}$$

$$\begin{aligned} &\mathrm{Tl}_2 \mathrm{O} \text { is as basic as } \mathrm{NaOH} \text { due to its lower oxidation state (+1). }\\ &\mathrm{Tl}_2 \mathrm{O}+2 \mathrm{HCl} \longrightarrow 2 \mathrm{TICl}+\mathrm{H}_2 \mathrm{O} \end{aligned}$$

P$$_{4}$$O$$_{10}$$ on reaction with water gives orthophosphoric acid

$$\underset{\begin{array}{c} \text { Phosphorus } \\ \text { pentoxide } \end{array}}{\mathrm{P}_4 \mathrm{O}_{10}}+6 \mathrm{H}_2 \mathrm{O} \longrightarrow \underset{\begin{array}{c} \text { Orthophosphoric } \\ \text { acid } \end{array}}{4 \mathrm{H}_3 \mathrm{PO}_4}$$

$$\begin{aligned} &\mathrm{Cl}_2 \mathrm{O}_7 \text { is strongly acidic in nature and on dissolution in water, it gives perchloric acid. }\\ &\underset{\substack{\text { Dichlorine heptoxide }}}{\mathrm{Cl}_2 \mathrm{O}_7}+\mathrm{H}_2 \mathrm{O} \longrightarrow \underset{\text { Perchloric acid }}{2 \mathrm{HClO}_4} \end{aligned}$$

First ionisation enthalpy of sodium $\left(\mathrm{Na}=1 s^2, 2 s^2, 2 p^6, 3 s^1\right)$ is lower than that of magnesium $\left(\mathrm{Mg}=1 s^2, 2 s^2 2 p^6, 3 s^2\right)$ because the electron to be removed in both the cases is from 3 s -orbital but the nuclear charge is lower in Na than that of magnesium. After the removal of first electron $\mathrm{Na}^{+}$acquires inert gas $(\mathrm{Ne})$ configuration $\left(\mathrm{Na}^{+}=1 s^2, 2 s^2, 2 p^6\right)$ and hence, removal of second electron from sodium is difficult. While in case of magnesium, after the removal of first electron, the electronic configuration of $\mathrm{Mg}^{+}$is $1 s^2, 2 s^2, 2 p^6, 3 s^1$. In this case $3 s^1$ electron is easy to remove in comparison to remove an electron from inert gas configuration. Therefore, $\mathrm{IE}_2$ of Na is higher than that of magnesium.

Exothermic reactions Reactions which are accompanied by evolution of heat are called exothermic reactions. The quantity of heat produced is shown either along with the products with a ' + ' sign or in terms if $\Delta H$ with a ' - ' sign. e.g.,

$$\begin{aligned} \mathrm{C}(\mathrm{s})+\mathrm{O}_2(g) & \longrightarrow \mathrm{CO}_2(g)+393.5 \mathrm{~kJ} \\ \mathrm{H}_2(g)+\frac{1}{2} \mathrm{O}_2(g) & \longrightarrow \mathrm{H}_2 \mathrm{O}(l) ; \Delta H=-285.8 \mathrm{~kJ} \mathrm{~mol}^{-1} \end{aligned}$$

Endothermic reactions Reactions which proceed with absorption of heat are called endothermic reactions. The quantity of heat absorbed is shown either along with the products with a ' - ' sign or in terms of $\Delta H$ with a ' + ' sign e.g.,

$$\begin{gathered} \mathrm{C}(\mathrm{s})+\mathrm{H}_2 \mathrm{O}(\mathrm{g}) \longrightarrow \mathrm{CO}(\mathrm{g})+\mathrm{H}_2(g)-131.4 \mathrm{~kJ} \\ \mathrm{~N}_2(g)+3 \mathrm{H}_2(g) \longrightarrow 2 \mathrm{NH}_3(g) ; \Delta H=+92.4 \mathrm{~kJ} \mathrm{~mol}^{-1} \end{gathered}$$

The placing of elements are as

(i) Ionisation enthalpy of nitrogen $\left({ }_7 \mathrm{~N}=1 s^2, 2 s^2, 2 p^3\right)$ is greater than oxygen $\left({ }_8 \mathrm{O}=1 s^2\right.$, $2 s^2, 2 p^4$ ) due to extra stable exactly half-filled $2 p$-orbitals. Similarly, ionisation enthalpy of phosphorus $\left({ }_{15} \mathrm{P}=1 s^2, 2 s^2, 2 p^6, 3 s^2, 3 p^3\right)$ is greater than sulphur $\left({ }_{16} \mathrm{~S}=1 s^2, 2 s^2, 2 p^6, 3 s^2, 3 p^4\right)$.

On moving down the group, ionisation enthalpy decreases with increasing atomic size. So, the order is

$$\mathrm{S}<\mathrm{P}<\mathrm{O}<\mathrm{N} \rightarrow \text { First ionisation enthalpy increases. }$$

(ii) Non-metallic character across a period (left to right) increases but on moving down the group it decreases. So, the order is

$\mathrm{P}<\mathrm{S}<\mathrm{N}<\mathrm{O} \rightarrow$ Non-metallic character increases.