Chemistry Group-2: Alkaline earth metals

### Topics to be covered

=> Alkaline earth metals
=> Electronic configuration
=> Atomic and ionic radii
=> Ionization enthalpies
=> Hydration enthalpies
=> Physical properties
=> Chemical properties

### GROUP 2 ELEMENTS : ALKALINE EARTH METALS

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.

### Electronic Configuration

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.

### Ionization Enthalpies

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.

### Hydration Enthalpies

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.

### Physical Properties

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.

### Chemical Properties

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.