`color{green}(★)` The group 2 elements comprise beryllium, magnesium, calcium, strontium, barium and radium.

`color{green}(★)` They follow alkali metals in the periodic table.

`color{green}(★)` These (except beryllium) are known as alkaline earth metals.

`color{green}(★)` The first element beryllium differs from the rest of the members and shows diagonal relationship to aluminium.

`color{green}(★)` The atomic and physical properties of the alkaline earth metals are shown in Table 10.2.

`color{green}(★)` These elements have two electrons in the `color{red}(s)` -orbital of the valence shell (Table 10.2).

`color{green}(★)` Their general electronic configuration may be represented as [noble gas] `color{red}(ns^2)`. Like alkali metals, the compounds of these elements are also predominantly ionic.

`color{green}( ★)` The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same periods.

`color{green}( •)` This is due to the increased nuclear charge in these elements. Within the group, the atomic and ionic radii increase with increase in atomic number.

`color{green}( ★)` The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms.

`color{green}( ★)` Since the atomic size increases down the group, their ionization enthalpy decreases (Table 10.2).


`color{green}( ★)` The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1 metals.

`color{green}( •)` This is due to their small size as compared to the corresponding alkali metals.

`color{green}( •)` It is interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.

`color{green}( ★)` Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group.

`color{red}(Be^(2+) > Mg^(2+) > Ca^(2+) > Sr^(2+) > Ba^(2+))`

`color{green}( ★)` The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.

`color{green}( ★)` Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, e.g., `color{red}(MgCl_2)` and `color{red}(CaCl_2)` exist as `color{red}(MgCl_2.6H_2O)` and `color{red}(CaCl_2· 6H_2O)` while `color{red}(NaCl)` and `color{red}(KCl)` do not form such hydrates.

`color{green}( ★)` The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals.

`color{green}( ★)` Beryllium and magnesium appear to be somewhat greyish.

`color{green}( ★)` The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes.

`color{green}( ★)` The trend is, however, not systematic. Because of the low ionisation enthalpies, they are strongly electropositive in nature.

`color{green}( ★)` The electropositive character increases down the group from `color{red}(Be)` to `color{red}(Ba.)`

`color{green}( ★)` Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours respectively to the flame.

`color{green}( ★)` In flame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light.

`color{green}( ★)` The electrons in beryllium and magnesium are too strongly bound to get excited by flame.

`color{green}( ★)` Hence, these elements do not impart any colour to the flame.

`color{green}( ★)` The flame test for `color{red}(Ca, Sr)` and `color{red}(Ba )` is helpful in their detection in qualitative analysis and estimation by flame photometry.

`color{green}( ★)` The alkaline earth metals like those of alkali metals have high electrical and thermal conductivities which are typical characteristics of metals.

`color{green}( ★)` The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.

`color{green}( ★ \ \ "𝐑𝐞𝐚𝐜𝐭𝐢𝐯𝐢𝐭𝐲 𝐭𝐨𝐰𝐚𝐫𝐝𝐬 𝐚𝐢𝐫 𝐚𝐧𝐝 𝐰𝐚𝐭𝐞𝐫 :")`

`color{green}( ★)` Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.

`color{green}( ★)` However, powdered beryllium burns brilliantly on ignition in air to give `color{red}(BeO)` and `color{red}(Be_3N_2)`.

`color{green}( ★)` Magnesium is more electropositive and burns with dazzling brilliance in air to give `color{red}(MgO)` and `color{red}(Mg_3N_2)`.

`color{green}( ★)` Calcium, strontium and barium are readily attacked by air to form the oxide and nitride.

`color{green}( ★)` They also react with water with increasing vigour even in cold to form hydroxides.

`color{green}( ★ \ \ "𝐑𝐞𝐚𝐜𝐭𝐢𝐯𝐢𝐭𝐲 𝐭𝐨𝐰𝐚𝐫𝐝𝐬 𝐭𝐡𝐞 𝐡𝐚𝐥𝐨𝐠𝐞𝐧𝐬 :")`

`color{green}( ★)` All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.

`color{red}(M+X_2 → MX_2 ( X = F , Cl , Br , I))`

`color{green}( ★)` Thermal decomposition of `color{red}((NH_4)_2BeF_4)` is the best route for the preparation of `color{red}(BeF_2)`, and

`color{red}(BeCl_2)` is conveniently made from the oxide.

`color{red}(BeO + C + Cl_2 overset(600-800K)⇌ BeCl_2 + CO)`

`color{green}( ★ \ \ "𝐑𝐞𝐚𝐜𝐭𝐢𝐯𝐢𝐭𝐲 𝐭𝐨𝐰𝐚𝐫𝐝𝐬 𝐡𝐲𝐝𝐫𝐨𝐠𝐞𝐧 :")`


`color{green}( ★)` All the elements except beryllium combine with hydrogen upon heating to form their hydrides,
`color{red}(MH_2).`

`color{green}( ★)` `color{red}(BeH_2)` however, can be prepared by the reaction of `color{red}(BeCl_2)` with `color{red}(LiAlH_4)`


`color{red}(2BeCl_2 + LiAlH_4 → 2 BeH_2 + LiCl + AlCl_3)`

`color{green}( ★ \ \ "𝐑𝐞𝐚𝐜𝐭𝐢𝐯𝐢𝐭𝐲 𝐭𝐨𝐰𝐚𝐫𝐝𝐬 𝐚𝐜𝐢𝐝𝐬 :")`

`color{green}( ★)` The alkaline earth metals readily react with acids liberating dihydrogen.

`color{red}(M + 2HCl → MCl_2 + H_2)`

`color{green}( ★ \ \ "𝐑𝐞𝐝𝐮𝐜𝐢𝐧𝐠 𝐧𝐚𝐭𝐮𝐫𝐞 :")`

`color{green}( ★)` Like alkali metals, the alkaline earth metals are strong reducing agents.

`color{green}( ★)` This is indicated by large negative values of their reduction potentials (Table 10.2).


`color{green}( ★)` However their reducing power is less than those of their corresponding alkali metals.

`color{green}( ★)` Beryllium has less negative value compared to other alkaline earth metals.

`color{green}( ★)` However, its reducing nature is due to large hydration energy associated with the small size of `color{red}(Be^(2+))` ion and relatively large value of the atomization enthalpy of the metal.

`color{green}( ★ \ \ "𝐒𝐨𝐥𝐮𝐭𝐢𝐨𝐧𝐬 𝐢𝐧 𝐥𝐢𝐪𝐮𝐢𝐝 𝐚𝐦𝐦𝐨𝐧𝐢𝐚 :")`

`color{green}( ★)` Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.

`color{red}(M + (x+y) NH_3 → [ M (NH_3)_X]^(2+) + 2 [ e (NH_3)_Y]^(-))`

`color{green}( ★)` From these solutions, the ammoniates, `color{red}([M(NH_3)_6]^(2+))` can be recovered.

`color{green}( ★ \ \ "𝐔𝐬𝐞𝐬 :")`

`color{green}( • )` Beryllium is used in the manufacture of alloys.

`color{green}( •)` Copper-beryllium alloys are used in the preparation of high strength springs.

`color{green}( • )` Metallic beryllium is used for making windows of X-ray tubes.

`color{green}( •)` Magnesium forms alloys with aluminium, zinc, manganese and tin.

`color{green}(•)` Magnesium-aluminium alloys being light in mass are used in air-craft construction.

`color{green}( •)` Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals.

`color{green}( •)` A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine.

`color{green}( •)` Magnesium carbonate is an ingredient of toothpaste.

`color{green}( •)` Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon.

`color{green}( •)` Calcium and barium metals, owing to their reactivity with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes.

`color{green}( •)` Radium salts are used in radiotherapy, for example, in the treatment of cancer.