Alkaline Earth Metals

Alkaline earth metals are a group of chemical elements found in Group 2 of the periodic table.

The alkaline earth family ranks as the second most reactive group among elements, and they're notably absent in their pure form in nature. They're labelled "alkaline" earth metals due to their tendency to produce "alkaline" solutions, specifically hydroxides, upon interaction with water. Interestingly, "earth" was the term alchemists used to describe the oxides derived from these alkaline earth metals.

1.0Definition of Alkaline Earth Metals

Elements within the alkaline earth metals exhibit an electronic configuration where their s-subshell holds two valence electrons. This configuration, [Noble gas] ns², characterises these metals. Positioned in the second column of the periodic table, they're commonly referred to as group two metals. For a detailed explanation, we will try to understand with examples of alkaline earth metals.

Alkaline earth metals cover a range of elements, such as Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). These elements share similar properties due to their two valence electrons and play vital roles in various industrial, biological, and chemical processes.

Furthermore, this section aims to cover some interesting and important alkaline earth metals characteristics.

2.0Physical Properties of Alkaline Earth Metals

Alkaline earth metals exhibit several key physical properties that distinguish them within the periodic table :

• Shiny Appearance: They typically have a shiny, metallic lustre when freshly cut.
• Malleability and Ductility: They are generally malleable, meaning they can be rolled into thin sheets, and ductile, allowing them to be drawn into wires.
• Good Conductors: These metals are efficient conductors of electricity and heat due to the mobility of their electrons.
• Low Density: Alkaline earth metals are relatively light in comparison to many other metals.
• Melting and Boiling Points: They possess moderately high melting and boiling points, with values that generally increase down the group.
• Reactivity: While less reactive than alkali metals, they are still notably reactive, particularly with water and oxygen in the air. They tarnish easily when exposed to air due to the formation of oxide layers on their surfaces.
• Atomic Size: They have relatively large atomic sizes within the periodic table, which increases down the group.
• Density: The density of alkaline earth metals generally increases down the group, with exceptions like magnesium, which is less dense than beryllium.

These physical properties collectively define the behaviour and characteristics of alkaline earth metals and influence their applications in various industries, including construction, metallurgy, medicine, and electronics.

3.0Chemical Properties of Alkaline Earth Metals

Here, we will discuss Alkaline earth metals chemical properties. Which also includes alkaline earth metals uses-

Reactivity: 

Alkaline earth metals are less reactive than alkali metals but still relatively reactive compared to other metals. They readily lose their two valence electrons to form ions with a +2 charge, becoming bivalent cations.

Formation of Oxides and Hydroxides: 

They react vigorously with oxygen to form oxides and hydroxides. For example, magnesium burns in the air to produce magnesium oxide. When these oxides dissolve in water, they form alkaline solutions (hence the name "alkaline earth metals").

Reducing Agents: 

They serve as reducing agents due to their ability to donate electrons. This property is particularly evident in reactions where they displace hydrogen from water or react with acids to produce hydrogen gas.

Formation of Salts: 

Alkaline earth metals form various salts, particularly sulphates, carbonates, and nitrates. These salts are widely used in industries such as agriculture, medicine, and manufacturing.

Alloys: 

They form alloys with other metals, enhancing the properties of the resulting material. For instance, magnesium alloys are used in aerospace and automotive industries due to their lightweight and high strength.

Coordination Chemistry: 

Alkaline earth metals play a role in coordination chemistry, often forming coordination complexes due to their ability to act as Lewis acids (electron-pair acceptors).

Reactivity with Halogens: 

They react with halogens to form ionic halides. These reactions are typically vigorous, especially with elements like chlorine or iodine.

4.0Radii of Alkaline Earth Metals

The atomic and ionic radii of the elements in the second group (alkaline earth metals) generally increase as you move down the group in the periodic table. Here's a brief overview:

• Atomic Radii: 

This refers to the size of the atom. As you move down the group from beryllium (Be) to radium (Ra), the atomic radius increases. This increase occurs because each successive element adds a new energy level (shell) to the electron cloud, increasing the distance between the nucleus and the outermost electrons.

• Ionic Radii: 

When alkaline earth metals lose their two valence electrons to form +2 cations (M²⁺), the ionic radius also follows a similar trend to the atomic radius. As you move down the group, the ionic radius of the +2 cations increases due to the addition of electron shells and increased electron-electron repulsion.

5.0Reactivity of Alkaline Earth Metals

Alkaline earth metals are relatively reactive, but their reactivity is lower compared to alkali metals in Group 1 of the periodic table. Let’s discuss their reactivity:

Reaction with Water: 

Alkaline earth metals react with water and give H₂, but their reactivity in water is less intense than alkali metals. For example:

• Beryllium doesn’t react with water due to the formation of a protective oxide layer on its surface.
• Magnesium reacts slowly with hot water, producing magnesium hydroxide and hydrogen gas.
• Calcium, strontium, barium, and radium react more vigorously with cold water, producing their respective hydroxides and hydrogen gas.

M   +   2 H₂O  →  M(OH)₂   +   H₂

Reaction with Oxygen: 

These metals readily react with oxygen in the air to form oxides. For example:

• Magnesium burns brightly in air to produce magnesium oxide.
• Calcium, strontium, barium, and radium also form oxides when exposed to air.

 MgO →  Weak basic

 CaO, SrO & BaO  →  Strong basic

• Basic properties increase from BeO to BaO.

 Reaction with Halogen- 

   Alkaline metals react with X (Halogen) to form MX₂.

• Order of Ionic nature

 BeCl₂   <   MgCl₂   <   CaCl₂ <   SrCl₂   <    BaCl₂

• The Hydrolysis tendency of these halides decreases from BeCl₂ to BaCl₂ due to decrease in covalent nature.
• BeCl₂ and MgCl₂ are covalent in nature.

BeO + C + Cl₂   →  BeCl₂ + CO    

Reaction with Hydrogen:

Except Be all the alkaline metals forms MH₂ type hydrides, (MgH₂, CaH₂, SrH₂, BaH₂) on heating directly with H₂

M + H₂   →  MH₂  

BeH₂ is prepared by action of BeCl₂ with  LiAlH₄ in presence of reducing agent  

2BeCl₂  + LiAlH₄   →  2BeH₂  + LiCl + AlCl₃

• Reaction with Acids: Alkaline earth metals react with acids, displacing hydrogen and forming salts. The reactivity generally increases down the group.

M + 2HCl → MCl₂ + H₂

• Formation of Salts: They form various salts, particularly sulphates, carbonates, and nitrates.
• Formation of Hydroxides: When these metals react with water, they form alkaline solutions due to the formation of hydroxides.

6.0Important Compounds of Alkaline Earth Metals

Some important compounds of the second group- Below we have covered a few important compounds and uses of alkaline earth metals. 

Calcium Oxide [Quick lime (CaO)]

Calcium oxide, commonly known as quicklime or burnt lime, is a white, caustic, alkaline substance.

Preparation of CaO :    

By heating limestone at high temperature (approx. 800⁰C-1000⁰C). 

CaCO₃  →   CaO  + CO₂

Properties  :

(i) Action of water :    

$\underset{\text{ (quick lime) }}{\mathrm{CaO}}+{\mathrm{H}}_2\mathrm{O}\to \underset{\text{ (Slaked lime) }}{\mathrm{Ca}(\mathrm{OH}{)}_2}$

Ca(OH)₂ Solution in water is called lime water.Ca(OH)₂ Suspension in water is called milk of lime

(ii) Basic Nature :

$\mathrm{CaO}+{\mathrm{SiO}}_2\underset{\text{ (Calcium silicate) }}{\overset{\Delta }{\to }}{\mathrm{CaSiO}}_3$

$\mathrm{CaO}+{\mathrm{P}}_4{\mathrm{O}}_{10}\stackrel{\Delta }{\to }2{\mathrm{Ca}}_3{\left({\mathrm{PO}}_4\right)}_2\text{ (Calcium phosphate) }$

(iii) Reaction with carbon:

$\mathrm{CaO}+3\mathrm{C}\stackrel{2000^{\circ }\mathrm{C}}{\to }\underset{\text{ (Calcium carbide) }}{{\mathrm{CaC}}_2+\mathrm{CO}\uparrow }$

${\mathrm{CaC}}_2+{\mathrm{N}}_2⟶\underset{\begin{array}{c}\text{ Nitrolime }\end{array}}{{\mathrm{CaCN}}_2+\mathrm{C}}$

Uses of CaO :

(i) In the manufacture of bleaching powder, cement, glass, calcium carbide etc.

(ii) In the purification of sugar            

(iii) As a drying agent for NH₃ and C₂H₅OH

(iv) As basic lining in furnaces

(v) For making Soda lime [NaOH + CaO]

Calcium Hydroxide [Slaked lime Ca(OH)₂]

Calcium hydroxide, commonly known as slaked lime or hydrated lime, is a chemical compound with the chemical formula Ca(OH)₂. 

Preparation of Ca(OH)₂ :  

By the action of water on quick lime

CaO   +   H₂O  →    Ca(OH)₂   +   heat (slaking of lime)

Properties of Ca(OH)₂

(i)       Action of CO₂ : Lime water turns milky on passing CO₂ gas. 

$\begin{array}{rl}& \mathrm{Ca}(\mathrm{OH}{)}_2+{\mathrm{CO}}_2⟶\underset{\text{ Milkiness }}{{\mathrm{CaCO}}_3}+{\mathrm{H}}_2\mathrm{O}\\ & {\mathrm{CaCO}}_3\underset{{\mathrm{CO}}_2+{\mathrm{H}}_2\mathrm{O}}{\overset{\text{ Excess of }}{\to }}\underset{\text{ (soluble) }}{\mathrm{Ca}{\left({\mathrm{HCO}}_3\right)}_2}\stackrel{\Delta }{\to }{\mathrm{CaCO}}_3\end{array}$

 (ii)     Action of Chlorine :

$2\mathrm{Ca}(\mathrm{OH}{)}_2+2{\mathrm{Cl}}_2\to \underset{\text{ Bleaching powder }}{{\mathrm{CaCl}}_2+\mathrm{Ca}(\mathrm{OCl}{)}_2}+2{\mathrm{H}}_2\mathrm{O}$

Uses of Ca(OH)₂-

(i) For softening of hard water.

(ii) For purification of sugar and Coal gas.

(iii) In the manufacture of bleaching powder, Caustic soda and soda lime

(iv) In preparation of mortar, plaster and white washing.

Calcium Sulphate [Gypsum CaSO₄.2H₂O]

Calcium sulphate dihydrate, commonly known as gypsum (CaSO₄·2H₂O), is a naturally occurring mineral compound. It appears as a soft, white or grey mineral found in sedimentary rock formations. Here are some key points about gypsum:

Preparation of Calcium Sulphate: CaSO₄.2H₂O is naturally occurring calcium sulphate. It can be obtained by the action of dilute H₂SO₄ on a soluble calcium salt below 60⁰C.

${\mathrm{CaCl}}_2+\underset{\text{ dilute }}{{\mathrm{H}}_2{\mathrm{SO}}_4}\to \underset{\text{ white ppt. }}{2\mathrm{HCl}}+{\mathrm{CaSO}}_4\downarrow$

Properties of Gypsum-

(i) Action of heat :

$\underset{\text{ (Gypsum) }}{2\left({\mathrm{CaSO}}_4.2{\mathrm{H}}_2\mathrm{O}\right)}\underset{-3{\mathrm{H}}_2\mathrm{O}}{\overset{120{}^{\circ }}{\to }}\underset{\text{ (Plaster of paris) }}{2\left({\mathrm{CaSO}}_4\right)\cdot {\mathrm{H}}_2\mathrm{O}}\stackrel{200^{\circ }\mathrm{C}}{\to }\underset{\text{ (Anhydride) }}{2{\mathrm{CaSO}}_4+{\mathrm{H}}_2\mathrm{O}}$

(ii) It forms an important fertiliser (NH₄)₂SO₄

CaSO₄ + 2NH₃ + CO₂ + H₂O   →    CaCO₃¯  + (NH₄)₂ SO₄

Uses of Gypsum-

(i) In the preparation of plaster of paris.

(ii) Anhydrous CaSO₄ used as a drying agent.

(iii) Anhydrite (CaSO₄) is used for manufacture of sulphuric acid, ammonium sulphate.