B.sc 1st year Book
(Page 6)

Group II A (2) Elements: Be, Mg, Ca, Sr, Ba, Ra (Alkaline Earth Metals)

The elements of group IIA of the periodic table consists of beryllium (Be) magnesium (Mg), calcium (Ca), strontium (SD), barium (Ba) and radium (Ra). These elements are commonly called ‘alkaline earth metals’ because they occur in the earth’s crust as oxides and their hydroxides behave as a strong base. Radium is a radioactive element but shows similar properties to alkaline earth metals hence placed in the same group.

1-Electronic configuration:

The general electronic configuration of IIA group elements is ns2,

2- Physical state :

All the elements of alkaline earth metals are soft with greyish-white lustre and were freshly cut. They are electropositive but less than group lA elements. They are malleable and ductile and harder than alkali metals.

3- Density :

Initially, the density of alkaline earth metals decreases slightly from Be to Ca and then there is a considerable increase in it from Ca to Ra as shown in the table.4.4. The value of densities of these metals are higher than alkali metals hence these metals are denser than ∣A group metals. This is due to the fact that the atoms in these elements are tightly packed because of an increase in the nuclear charge and they have a smaller size as compared to alkali metals.

4- Melting and boiling points :

There is no regular trend of variation in m.p. and b.p. of alkaline earth metals as we move from top to bottom in a group. But the values of m.p. and b.p. are higher than the IA group elements. Amongst them, Be shows a very high value of m,p. and b,p. which is due to the small. size of atoms which are tightly packed.

5 – Atomic and ionic radii :

(Atomic volume) In a group, or moving from top to bottom the values of atomic and ionic radii increases with an increase in atomic number. This increase is due to an increase in the number of shells. These values of atomic and ionic radii are lower than the respective values of alkali metals. The size of the atom and Ion decreases with an increase in nuclear charge which attracts electrons towards itself as compared to lA group elements. Due to smaller atomic radii, these metals are harder and have higher values of density and m.p. and b.p. as shown in table 4.4.

6- Colour and magnetic properties :

All the alkaline earth metals form M²⁺ ions whose electronic configuration is similar to inert gas like 1s² or ns²np⁶. These Ions have no unpaired electron hence they show diamagnetism and are colourless.

7-Ionization Energy:

The values of first and second ionization energies of these elements decrease from top to bottom in a group, Ra has a slightly greater value of I.E. as compared to Ba. The value of I.E. of ll group elements is greater than that of alkali metals. This is due to their small size and greater nuclear charge. These metals have a strong electropositive character which increases from Be to Ba but they are less electropositive than alkali metals. Yet the value of ionization energy of these elements is low when they form M⁺ ions after easily losing one electron from their outermost shell (ns²). But these metals prefer to form M²⁺ ions to get the inert gas configuration, (like 1s² or ns²np⁶). They never form M³⁺ ions because the removal of electrons from inert gas configuration is very difficult.

8- Metallic and electropositive character :

The ionization energies of alkaline earth metals are low. Hence they are highly electropositive but less than that of alkali metals. Unlike alkali metals, alkaline earth metals do not emit electrons on exposure to light. Since I.E. decreases from Be to Ba the electropositive character increases. Due to this, the basic character of hydroxides increases from Be to Ba. Be(OH )₂ behaves as an amphoteric hydroxide because it reacts with acids and bases both to form salt and Beryilates (BeO²₂₋).

9. Oxidation states :

Since all the alkali metals have two-electron in their outermost orbit hence they show +2 oxidation states. But in terms of ionization energy one can easily predict the existence of univalent metal ions. This anomaly can be explained on the basis of solvation energy or hydration energy.

10- Hydration of M²⁺ ions :

In IIA group the size of M²⁺ ions increases from top to bottom causing a decrease in the value of hydration energy. When we compare the value of hydration energies of M²+1 in of Ingroup and the M⁺ ion of the IIA group, It was found that M²⁺ ions are more extensively hydrated than those of alkali metals. The decrease in the values of hydration energies from Be²⁺ to Ba²⁺ increases the ionic mobility of ionic conductance of M2+ cations.

11- Reducing property :

Since all the alkaline earth metals have a tendency to lose electrons and form M²⁺ ions hence they act as strong reducing agents but the reducing property of these metal ions is weaker than alkali metal ions. In a group the reducing property increases from Be to Ba regularly as shown below:

Since Be has the lowest value of Ee hence it liberates H2 from dilute acids very slowly as compared to other members and is the weakest reducing agent.

12- Electronegativity :

Since all the alkaline earth metal ions are highly electropositive. Hence they have low electronegativity values like alkali metals, The electronegativity values decrease from Be to Ba with the increase in the size of atoms.

13- Colour of flame :

Amongst alkaline earth metals Be and Mg do not give any flame colour whereas Ca, Sr, Be and Ra produce different colours when subjected to a flame test. The colour of the flame is depicted in table 4.4. Colour is produced when metal or metal salts are heated in a Bunsen flame. The electron absorbs energy and becomes excited to a higher energy level. After returning to the ground state from the excited state absorbed energy is released in the form of visible light which imparts characteristic flame colour. Due to the small size of Be and Mg, they do not impart much colour because electrons are held up by the nucleus tightly and the energy released is in the u.v. region.

   Page No. 6

   Previous PageNext Page