The reactivity of alkali metals towards oxygen increases on moving down the group with the increase in atomic size. Thus, Li forms only lithium oxide $\left(\mathrm{Li}_2 \mathrm{O}\right)$, sodium forms mainly sodium peroxide $\mathrm{Na}_2 \mathrm{O}_2$ along with a small amount of sodium oxide while potassium forms only potassium superoxide $\left(\mathrm{KO}_2\right)$.

$4 \mathrm{Li}+\mathrm{O}_2 \xrightarrow{\Delta} 2 \mathrm{Li}_2 \mathrm{O}$

$6 \mathrm{Na}+2 \mathrm{O}_2 \xrightarrow{\Delta} \underset{\substack{\text { Sodium peroxide } \\ \text { (major) }}}{\mathrm{Na}_2 \mathrm{O}_2}+\underset{\substack{\text { Monooxide } \\ \text { (minor) }}}{2 \mathrm{Na}_2 \mathrm{O}}$

$\mathrm{K}+\mathrm{O}_2 \xrightarrow{\Delta} \underset{\begin{array}{c}\text { Potassium } \\ \text { super oxides }\end{array}}{\mathrm{KO}_2}+\underset{\text { Peroxide }}{\mathrm{K}_2 \mathrm{O}_2}+\underset{\text { Monoxide }}{\mathrm{K}_2(\mathrm{O})}$

The superoxide, $\mathrm{O}_2^{-}$ion is stable only in presence of large cations such as $\mathrm{K}, \mathrm{Rb}$ etc.

$\mathrm{O}_2{ }^{2-}$ represents a peroxide ion

$\mathrm{O}_2^-$ represents a superoxide ion

(i) Peroxide ion react with water to form $\mathrm{H}_2 \mathrm{O}_2$

$$\mathrm{O_2^{2 - } + 2{H_2}O\buildrel {} \over \longrightarrow 2O{H^ - } + \mathop {{H_2}{O_2}}\limits_{Hydrogen\,peroxide}}$$

(ii) Superoxide ion react with water to form $\mathrm{H}_2 \mathrm{O}_2$ and $\mathrm{O}_2$

$2 \mathrm{O}_2^{-}+2 \mathrm{H}_2 \mathrm{O} \longrightarrow 2 \mathrm{OH}^{-}+\underset{\substack{\text { Hydrogen } \\ \text { peroxide }}}{\mathrm{H}_2 \mathrm{O}_2}+\mathrm{O}_2$

Lithium resembles with magnesium as its charge size ratio is closer to Mg. Its resemblance with Mg is known as diagonal relationship.

Generally, the periodic properties show either increasing or decreasing trend along the group and vice-versa along the period which brought the diagonally situated elements to closer value.

Following characteristics can be noted

(i) Due to covalent nature, chlorides of Li and Mg are deliquescent and soluble in alcohol and pyridine.

(ii) Carbonates of Li and Mg decompose on heating and liberate $\mathrm{CO}_2$

$\begin{aligned} \mathrm{Li}_2 \mathrm{CO}_3 & \longrightarrow \mathrm{Li}_2 \mathrm{O}+\mathrm{CO}_2 \\ \mathrm{MgCO}_3 & \longrightarrow \mathrm{MgO}+\mathrm{CO}_2\end{aligned}$

An element from group 2 which forms an amphoteric oxide and a water soluble sulphate is beryllium.

Beryllium forms oxides of formula BeO . All other alkaline earth metal oxides are basic in nature. BeO is amphoteric in nature i.e., it reacts with acids and bases both.

$$\begin{gathered} \mathrm{Al}_2 \mathrm{O}_3+2 \mathrm{NaOH} \longrightarrow 2 \mathrm{NaAlO}_2+\mathrm{H}_2 \mathrm{O} \\ \mathrm{Al}_2 \mathrm{O}_3+6 \mathrm{HCl} \longrightarrow 2 \mathrm{AlCl}_3+3 \mathrm{H}_2 \mathrm{O} \end{gathered}$$

Sulphate of beryllium is a white solid which crystallises as hydrated salts ( $\left.\mathrm{BeSO}_4 \cdot 4 \mathrm{H}_2 \mathrm{O}\right)$. $\mathrm{BeSO}_4$ is fairly soluble in water due to highest hydration energy in the group (small size). For $\mathrm{BeSO}_4$, hydration energy is more than lattice energy and so, they are readily soluble.

(i) All the alkaline earth melals form carbonates $\left(\mathrm{MCO}_3\right)$. All these carbonates decompose on heating to give $\mathrm{CO}_2$ and metal oxide. The thermal stability of these carbonates increases down the group i.e., from Be to Ba .

$$\mathrm{BeCO}_3<\mathrm{MgCO}_3<\mathrm{CaCO}_3<\mathrm{SCO}_3<\mathrm{BaCO}_3$$

$\mathrm{BeCO}_3$ is unstable to the extent that it is stable only in atmosphere of $\mathrm{CO}_2$. These carbonates however show reversible decomposition in closed container.

$$\mathrm{BeCO}_3 \rightleftharpoons \mathrm{BeO}+\mathrm{CO}_2$$

Hence, more is the stability of oxide formed, less will be stability of carbonates. Stability of oxides decreases down the group is beryllium oxide i.e., high stable making $\mathrm{BeCO}_3$ unstable.

(ii) All the alkaline earth metals form oxides of formula $M O$. The oxides are very stable due to high lattice energy and are used as refractory material.

Except BeO (predominantly covalent) all other oxides are ionic and their lattice energy decreases as the size of cation increases.

The oxides are basic and basic nature increases from BeO to BaO (due to increasing ionic nature).

$\underbrace{\mathrm{BeO}}_{\text {Amphoteric }}<\underbrace{\mathrm{MgO}<}_{\text {Weak basic }} \underbrace{\mathrm{CaO}<\mathrm{SrO}<\mathrm{BaO}}_{\text {Strong basic }}$

BeO dissolves both in acid and alkalies to give salts and is amphoteric The oxides of the alkaline earth metals (except BeO and MgO ) dissolve in water to form basic hydroxides and evolve a large amount of heat. BeO and MgO possess high lattice energy and thus insoluble in water.