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Alkaline Earth Metals
The group 2 of the periodic table consists of six metallic elements. They are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra). The name alkaline earth metals was given to magnesium, calcium, barium & strontium since their oxides were alkaline in nature and these oxide remained unaffected by heat or fire and existed in earth.
Like alkali metals, alkaline earth metals are also highly reactive and hence do not occur in the free state but are likely distributed in nature in the combined state as silicates, carbonates, sulphates and phosphates.
Be – Beryl (Be3Al2Si6O18) & Phenacite (Be2SiO4)
Mg – Magnesite MgCO3, Dolomite CaMg(CO3)2, Epsomite MgSO4.&H2O
Ca – Limestone (CaCO3), fluoropatite [3(Ca3(PO4)3.CaF2],
Gypsum (CaSO4.2H2O), Anhydrite (CaSO4) Sr – Celestite (SrSO4),
Br – Barytes (BaSO4)
The general electronic configuration of alkaline earth metals is ns2.
Be – 1s22s2 Mg – 1s22s2sp63s2
Ca – 1s22s22p63s23p64s2 Sr – [Kr]5s2
Ba – [Xe]6s2 Ra – [Rn]7s2
Physical Properties of Group II elements
Atomic and ionic radii
The atomic radii as well as ionic radii of the members of the family are smaller than the corresponding members of alkali metals.
The alkaline earth metals owing to their large size of atoms have fairly low values of ionization energies as compared to the p – block elements. However with in the group, the ionization energy decreases as the atomic number increases. It is because of increase in atomic size due to addition of new shells and increase in the magnitude of screening effect of the electrons in inner shells. Because their (IE)1 is larger than that of their alkali metal neighbours, the group IIA metals trend to the some what less reactive than alkali metals. The general reactivity trend is Ba > Sr > Ca > Mg > Be.
The 2nd ionization energies of the elements of group I are higher than those of the elements of group II. Explain.
The2nd electron in case of alkali metal is to be removed form a cation (unipostive ion) which has already acquired a noble gas configuration whereas in case of alkaline earth metals, the second electron is to be removes fro a cation which is yet to acquire the stable noble gas configuration therefore, removal of 2nd electron in case of alkaline earth metals requires much less energy than that in case of alkali metals.
There is sharp increase in third ionization energy due to stable inert gas configuration of m+2 ions. This explains the upper limit of +2 oxidation state for the elements.
The alkaline earth metal have two electrons in their valence shell and by losing these electrons, these atoms acquire the stable noble gas configuration. Thus, unlike alkali metals, the alkaline earth metals exhibit +2 oxidation state in their compounds.
M → M+2 + 2e-
The alkaline earth metals shows +2 oxidation state i.e. they always form divalent cations (M2+). Explain.
If ionization energy were the only factor involved, than group II elements should have formed monovalent ions i.e. Mg+, Ca+ etc rather than Mg2+, Ca+2 etc.
This can be explained as follows:
(i) The divalent cations of alkaline earth metals acquires stable inert gas configuration.
(ii) The divalent cations results in stronger lattices then monovalent cations and hence a lot of energy called lattice energy released during formation of divalent cations than monovalent cation which compensates the high second ionization energy.
(iii) The existence of divalent ions in the aqueous solution is due to greater hydration of the divalent ions which counter balance the high value of second ionization energy.
The heat of hydration (hydration energy) of alkaline earth metals are approximately four times higher than alkali metals of comparable size. e.g.
ΔHhyd for Na+ (size 102 pm) = -397 KJmol-1
ΔHhyd for Ca+2 (size 100 pm) = -1650 KJmol-1
Larger hydration energy is due to the fact that the alkaline earth metals ions, because of their much larger charge to size ratio, exert a much stronger electrostatic attraction on the oxygen of water molecule.