In group 2 of the periodic table, there are six elements. They are so named because the oxides that are found in nature are alkaline or basic. They remain unaffected by any combusting agent like heat or fire. They are highly reactive and do not exist in the earth’s crust in their free states. Rather, they are found in their combined states like silicates, carbonates, etc.

Occurrence on Earth

• Beryllium → is found in silicate minerals in small quantities.
• Magnesium → abundantly found as carbonates, silicates and silicates. It is the sixth most abundant element found on earth.
• Calcium → is the fifth most abundant element found in the earth’s crust. It is found in the form of calcium carbonate in limestone, chalks and marble. Other examples include gypsum, fluorspar, to name a few.
• Strontium → is mined in the form of strontianite (SrCO3) or celestite (SrSO4).
• Barium → they are mined in the form of barytes. 
• Radium → is a radioactive element. It is found in the form of disintegrated states of uranium, thorium and actinium.

Physical Properties

• Electronic configuration and Nature of bonds:

(noble gas) ns2 is the general electronic configuration. The atoms of these elements have a filled s-subshell. 

• Ionic and atomic radii:

Ionic and atomic radii keep on increasing as we go down the group in the periodic table. They can lose all of their electrons from the s-subshell, becoming smaller than the neutral atom in size. But even so, they still keep on increasing down the group.

• Ionisation and hydration energies/enthalpies:

They have low ionisation enthalpies but are greater than those of group 1 elements in the periodic table for the following reasons:

• Due to the high nuclear charge and small radii, they are held much more tightly.
• Because of the stable subshell, removing an electron requires a huge amount of energy.

The second ionisation enthalpies are low as compared to the group 1 elements. They have a half-filled s-subshell, and removing an electron is feasible by using less energy to attain the noble gas octet configuration.

Hydration enthalpy is inversely proportional to the atomic radii. As the size increases down the group, hydration enthalpy decreases.

• Boiling and melting points:

These elements have relatively high melting and boiling points compared to the alkali metals because of their small size and strong forces of attraction between them in the crystal lattices. Down the group, leaving magnesium, boiling, and melting points decrease.

• Electropositivity of alkaline earth metals:

The alkaline earth metals are highly electropositive or metallic, and their metallic character keeps on increasing as we go down the group. Due to low ionisation enthalpies, they tend to lose electrons, easily forming divalent cations increasing their metallic character. 

Alkaline earth metals have a lower metallic character than alkali metals. Because of their smaller size than alkali metals, they experience a greater nuclear charge reducing the tendency to lose electrons more easily than the group 1 elements.

• Nature of bonds:

As we go down the group, the tendency to form ionic bonds increases. As the ionisation enthalpy goes on decreasing down the group, the nature to form ionic bonds increases. Because of decreasing ionisation energy, the atoms can lose electrons completely and form ionic bonds. Beryllium and magnesium show some sort of covalency but the rest of all other bonds are ionic. 

• Density:

The density of these elements do not show any such trend down the group, but they are denser than the group 1 elements because of the small size and strong forces of attraction between them. However, there is a decrease in density from beryllium to calcium and increases from calcium to barium. The decrease in density can be explained due to the decrease in the packing efficiency of the atoms in the crystal lattice.

• Flame test colouration:

Just like group 1 elements, alkaline earth metals also exhibit and impart a characteristic flame to distinguish them from the rest. When these elements are heated to a higher level, the electrons in the atoms get excited and move to a higher state. Upon returning to the ground state, the absorbed energy is released in the form of light of certain wavelengths. 

• brick red colour imparted by calcium
• crimson red colour imparted by strontium
• apple green colour imparted by barium
• crimson colour imparted by radium

Being small in size, beryllium and magnesium do not impart any colour because they are held strongly by the nucleus. A burner does not provide the required heat for those electrons to leave the ground state. Hence, they do not lose electrons so easily. 

Physical properties of alkaline earth metals importance

Showing a high metallic character, alkaline earth metals are very good conductors of heat and electricity. They have a shiny lustre but get tarnished very readily when exposed to air. It is very hard to scratch glass with beryllium. On the contrary, barium is a bit harder than lead.

Conclusion

The alkaline earth metals are reactive electropositive or metallic elements that have a greyish-white, shiny lustre in them. They belong to the second group of the periodic table, having a filled s-subshell. The elements in this group have almost similar properties both physically and chemically. Owing to the physical properties of alkaline earth metals importance, they have been found very useful in different fields of application.