Physical properties

   (1) Physical state : All are greyish- light, white, malleable and ductile metals with metallic luster. Their hardness decrease progressively with increase in the atomic number. Though these are fairly soft but relatively harder than the alkali metals.

(2) Atomic and ionic radii

         (i) The atomic and ionic radii of alkaline earth metals also increase down the group due to progressive addition of new energy shells like alkali metals.

                                                     Be         Mg         Ca        Sr         Ba           Ra

Atomic radius (pm)             112        160        197      215        222          -

Ionic radius of M²⁺               31           65         99       113         135       140

ion (pm)

   (ii) The atomic radii of alkaline earth metals are however smaller than their corresponding alkali metal of the same period. This is because of the fact that alkaline earth metals possess a higher nuclear charge than alkali metals which more effectively pulls the orbit electrons towards the nucleus causing a decrease in size.

         (3) Density  

         (i) Density decreases slightly upto Ca after which it increases. Decrease in the density from Be to Ca may be due to less packing of atoms in solid lattice of Mg and Ca.

          Be          Mg       Ca       Sr        Ba      Ra

   1.84      1.74   1.55   2.54   3.75   6.00

(ii) The alkaline earth metals are more heavier, denser,  and harder than the alkali metal. The higher density of the alkaline earth metals is because of their smaller atomic size and strong intermetallic bonds which provide a more close packing in crystal lattice as compared to alkali metals.

         (4) Melting point and Boiling point

         (i) The melting points and boiling points of the alkaline earth metals do not show any regular trend.

                                                       Be         Mg        Ca          Sr           Ba            Ra

melting points (K)            1560     920      1112      1041      1000          973

boiling point (K)              2770    1378     1767      1654      1413           -

(ii) The values are, though, more than alkali metals. This can be due to close packing of atoms in crystal lattice in alkaline earth metals.

         (5) Ionisation energy and electropositive or metallic character

         (i) Since the atomic size decreases along the period and the nuclear charge increases and thus the electrons are more tightly held towards nucleus. It is hence alkaline earth metals have higher ionisation energy in comparison to the alkali metals but lower ionisation energies in comparison to the p-block elements. 

         (ii) The ionisation energy of the alkaline earth metals decreases from Be to Ba.

                                                                       Be        Mg       Ca       Sr         Ba         Ra

First ionisation energy (k J mol^-1)       899   737    590      549    503     509

Second ionisation  energy (kJ mol^-1) 1757  1450  1146   1064   965    979

(iii) The higher values of the second ionisation energy is because of the fact that removal of one electron from valence shell, the remaining electrons are more tightly held in which the nucleus of cation and hence more energy is needed to pull one more electron from the monovalent cation.

   (iv) No doubt first ionisation energy of alkaline earth metals are higher than alkali metals but a closer look on 2nd ionisation energy of alkali metals and alkaline earth metals reveals that 2nd ionisation energy of alkali metals are more

                                                                                 Li          Be

1st ionisation energy (kJ mol^-1)                      520    899  

2nd ionisation energy (kJ mol^-1)                 7296    1757

         This may be explained as,

         The removal of 2^nd electron from alkali metals takes place from 1s sub shell which are more closer to nucleus and exert more nuclear charge to hold up 1s electron core, while removal of 2nd  electron from alkaline earth metals takes from 2s sub shell. Closer are shells to the nucleus, more tightly will be held electrons with the nucleus and thus more energy is needed to remove the electron.

         (v) All these have strong electropositive character which increases from Be to Ba.

         (vi) These posses less electropositive character than alkali metals as the later have low values of the ionisation energy.

         (6) Oxidation number and valency  

         (i) The IE₁ of these metals are much lower than IE₁ and therefore it appears that these metals should form univalent ion rather than divalent ions but in actual practice, all these provide bivalent ions. This is because of the fact that M²⁺ ion possesses a higher degree of hydration or M²⁺ ions are extensively hydrated to form the [M(H₂O)_x]²⁺, a hydrated ion. This includes a large amount of energy evolution which counter balances the higher value of second ionisation energy.

            M → M²⁺ ,     DH = IE₁ + E₂

            M²⁺ + _xH₂O → [M(H₂O)_x]²⁺; DH = - hydration energy.

         (ii) The tendency of these metals to exist as divalent cation can hence be accounted as,

     (a) Divalent cation of these metals have noble gas or stable configuration.

         (b) Formation of the divalent cation lattice leads to the evolution of energy because of strong lattice structure of divalent cation which easily compensates for higher values of second ionisation energy of these metals.

         (c) Higher heats of the hydration of divalent cation which accounts for the existence of the divalent ions of these metals in solution state.

         (7) Hydration of ions

         (i) The hydration energies of alkaline earth metals divalent cation are much more than the hydration energy of monovalent cation.

                                                                                      Mg⁺       Mg²⁺

Hydration energy or Heat of hydration (kJ mol^-1)          353    1906

         The abnormally higher values of heat of hydration for divalent cations of alkaline earth metals are responsible for their divalent nature. MgCl₂ formation occurs with more amount of heat evolution and thus MgCl₂ is more stable.

         (ii) The hydration energies of M²⁺ ion decreases with increase in ionic radii.

                                                Be²⁺       Mg²⁺           Ca²⁺       Sr²⁺               Ba²⁺

Heat of hydration kJ mol^-1     2382     1906        1651     1484        1275

         (iii) Heat of the hydration are larger than the alkali metals ions and thus alkaline earth metals compounds are more extensively hydrated than those of alkali metals e.g MgCl₂ and CaCl₂ exists as Mg Cl₂ .6H₂O and CaCl_2. 6H₂O which NaCl and KCl do not form such hydrates.

         (iv) The ionic mobility, thus, increases from Be²⁺ to Ba²⁺, as the size of hydrated ion decreases.

         (8) Electronegativities

         (i) The electronegativities of alkaline earth metals are also small but are higher than alkali metals.

         (ii) Electronegativity decreases from Be to Ba as shown below,

                                     Be      Mg      Ca       Sr        Ba      

Electronegativity 1.57   1.31   1.00   0.95   0.89  

         (9) Conduction power : Good conductor of heat and electricity.

         (10) Standard oxidation potential and reducing properties

         (i) The standard oxidation potential (in volts) are,

        
              Be     Mg      Ca       Sr        Ba

            1.69   2.35   2.87   2.89   2.90

         (ii) All these metals possess tendency to lose two electrons to give M²⁺ ion and are used as reducing agent.

         (iii) The reducing character increases from Be to Ba, though, these are less powerful reducing agent than alkali metals.

         (iv) Beryllium possesing relatively lower oxidation potential and thus does not liberate H₂ from acids.
         (11)  Characteristic flame colours

         The characteristic flame colour shown are : Ca - brick red; Sr -crimson ; Ba-apple green and Ra- crimson.

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