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Chapter: 11th 12th std standard Class Organic Inorganic Physical Chemistry Higher secondary school College Notes

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Group 2 s - Block Elements

Group 2 s - Block Elements
The second group of the periodic table contains Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra). These elements are also a known as "Alkaline Earth Metals".



The second group of the periodic table contains Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra). These elements are also a known as "Alkaline Earth Metals". The word earth was applied in old days to a metallic oxide and because the oxides of calcium, strontium and barium produced alkaline solutions in water and, therefore these metals are called the alkaline earth metals. Radium corresponds to all the alkaline earth metals in its chemical properties but being radioactive, it is studied along with other radioactive elements.


Like the alkali metals, they are very reactive and hence never occur in nature in free form and react readily with many non metals.


Electronic configuration

Element      At      Electronic   Configuration of

         No.             Valence Shell


Beryllium   (4)  : 1s22s2        - 2s2

Magnesium (12) :           1s22s22p63s2      - 3s2

Calcium      (20) :           1s22s22p63s23p64s2    - 4s2

Strontium   (38) :           1s22s22p63s23p64s24p6 5s2 - 5s2

                   1s22s22p63s23p63d10 4s24p64d10         - 2

Barium       (56) :           5s2 5p66s2 6s

                   1s22s22p63s23p63d104s2 4p64d10         - 2

Radium       (88) :           5s2 5p6 5d10 5f14 6s2 6p67s2        - 7s

The electronic configurations show that for each element, the neutral atom has two electron after inert gas core and two electrons are in a completed s-subshell. Thus, the outer electronic configuration of each element is ns2 where n is the number of the valence shell. It can be expected that the two electrons can be easily removed to give the inert gas electronic configuration. Hence these elements are all bivalent and tend to form ionic salts. Thus ionic salts are less basic than group 1. Due to their alike electronic structure, these elements resemble closely in physical and chemical properties.


The variation in physical properties are not as regular as for the alkalimetals because the elements of this group do not crystallise with the same type of metallic lattice.


These elements have been sufficiently soft yet less than the alkalimetals as metallic bonding in these elements has been stronger than in first group alkali elements.


Beryllium is unfamiliar, partly because it is not very abundant and partly because it is difficult to extract. Magnesium and calcium are abundant and among the eight most common elements in the earth's curst. Strontium and barium are less abundant but are well known, while radium is extremely scarce and its radioactivity is more important than its chemistry.

Metallic properties


The alkaline earth metals are harder than the alkali metals. Hardness decreases, with increase in atomic number. They show good metallic lustre and high electrical as well as thermal conductivity because the two s-electrons can easily move through the crystal lattice.

Melting and Boiling Points


Both melting and boiling points do not show regular trends because atoms adopt different crystal structures. They possess low melting and boiling points. These are, however, higher than those of alkali metals because the number of bonding electrons in these elements is twice as great as group 1 elements.

Atomic radius


The atoms of these elements are somewhat smaller than the atoms of the corresponding alkali metals in the same period. This is due to higher nuclear charge of these atoms which tends to draw the orbital electrons inwards. Due to the smaller atomic radius, the elements, are harder, have higher melting points and higher densities than the elements of group 1. Atomic radius is seen to increase on moving down the group on account of the presence of an extra shell of electron at each step.

Ionic radius


The ions are also large but smaller than those of the elements in group 1. This is due to the fact that the removal of two orbital electrons in the formation of bivalent cations M2+, (Be2+, Mg2+, Ca2+, Sr2+, etc) increases the effective nuclear charge which pulls the electrons inwards and thus reduces the size of the ions. The ionic radius is seen to increase on moving down the group 2.

Atomic volume


Due to the addition of an extra shell of electrons to each element from Be to Ra, the atomic volume increases from Be to Ra.

Ionisation Energy


As the alkaline earth metals are having smaller size and greater nuclear charge than the alkali metals, the electrons are more tightly held and hence the first ionisation energy would be greater than that of the alkali metal.


The second ionisation energy has been to be nearly double than that of the first ionisation energy.


It is interesting to observe that although the IE2 of the alkaline earth metals is much higher than the IE1 they are able to form, M2+ ions. This is due to their high heat of hydration in aqueous solution and high lattice energy in the solid state. As the atomic size gets increased from Be to Ba, the values of IE1 and IE2 of these elements would decrease on going down the group, ie, Be to Ba.


As among second group elements beryllium has the highest ionisation energy. It has the least tendency to form Be2+ ion.


Thus its compounds with nitrogen, oxygen, sulphur and halogens are covalent whereas the corresponding compounds of Mg, Ca, Sr and Ba are ionic.

The total energy required to produce gaseous divalent ion for second group elements is over four times greater than the amount needed to ionise alkali metals. This very high energy requirement is more than offset by the hydration energy or the lattice energy being more than four times greater.

Oxidation states


Because of the presence of two s-electrons in the outermost orbital, being high heat of hydration of the dipositive ions and comparatively low value of IE2, the alkaline earth metals have been bivalent. The divalent ion is having no unpaired electron, hence their compounds are diamagnetic and colourless, provided their anions have been also colourless.

Flame colouration


These elements and their compounds impart characteristic colours to flame. Thus, barium - apple green, calcium - brick red, strontium - crimson red, radium - crimson red.


The reason for imparting the colour to flame is that when elements or their compounds are put into flame, the electrons get energy and excite to higher energy levels. When they return to the ground state they emit the absorbed energy in the form of radiations having particular wavelength.


Beryllium and magnesium atoms are smaller and their electrons being strongly bound to the nucleus are not excited to higher energy levels. Therefore they do not give the flame test.

Diagonal relationship between Beryllium and Aluminium


In case of beryllium, a member of second period of the periodic table, which resembles more with Aluminium group (13 group) than the member of its own group (2nd). The anamolous behaviour of beryllium is mainly ascribed to its very small size and partly due to its high electronegativity. These two factors tend to increase the polarising power of Be2+ tends to form ions to such extent that it is significantly equal to the polarising power of Al3+ ions. Thus the two elements resemble very much.



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