Explain the chemical properties of alkaline earth metals (Group $2$) in detail.

(N/A) $(i)$ Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.
However,powdered beryllium burns brilliantly on ignition in air to give $BeO$ and $Be_{3}N_{2}$.
Magnesium is more electropositive and burns with dazzling brilliance in air to give $MgO$ and $Mg_{3}N_{2}$.
Calcium,strontium,and barium are readily attacked by air to form the oxide and nitride. They also react with water with increasing vigour even in cold to form hydroxides.
$(ii)$ Reactivity towards the halogens: All the alkaline earth metals combine with halogens at elevated temperatures forming their halides,$MX_{2}$ $(X = F, Cl, Br, I)$.
Thermal decomposition of $(NH_{4})_{2}BeF_{4}$ is the best route for the preparation of $BeF_{2}$,and $BeCl_{2}$ is conveniently made from the oxide: $BeO + C + Cl_{2} \xrightarrow{600-800 \ K} BeCl_{2} + CO$.
$(iii)$ Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides,$MH_{2}$. $BeH_{2}$,however,can be prepared by the reaction of $BeCl_{2}$ with $LiAlH_{4}$: $2BeCl_{2} + LiAlH_{4} \rightarrow 2BeH_{2} + LiCl + AlCl_{3}$.
$(iv)$ Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen: $M + 2HCl \rightarrow MCl_{2} + H_{2}$.
$(v)$ Reducing nature: Like alkali metals,the alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials. However,their reducing power is less than those of their corresponding alkali metals.
$(vi)$ Solutions in liquid ammonia: Like alkali metals,the alkaline earth metals dissolve in liquid ammonia to give deep blue-black solutions forming ammoniated ions: $M_{(s)} + (x+y)NH_{3} \rightarrow [M(NH_{3})_{x}]^{2+} + 2[e(NH_{3})_{y}]^{-}$.