POTASSIUM (K)
Source:
Sylvine – KCl
Carnallite – KCl.MgCl2. 6H2O
Long benite – K2SO4. 2Mg SO4
Schonite – K2SO4. MgSO4. 6H2O
Poly halite – K2SO4. MgSO4. CaSO4. 2H2O
Kainite – KCl. MgSO4.3H2O
Salt petre – KNO3
Feldspar – KAlSi3O8
Extraction
(1) By heating K2CO3 with charcoal
K2CO3 + 2C → 2K + 3CO
(2) By reduction of potassium fluoride with calcium carbide
2KF + CaC2 → 2K + CaF2 + 2C.
Properties:
(1) It is silvery white metal
(2) Its all the properties are similar to that of sodium.
Group – 2: Alklaline Earth Metals
All alkaline earth metals have two s-electrons in their outermost shell over a noble gas core.
Electronic Configuration of Alkaline Earth Metals
_________________________________________________
Element with Configuration
At. No. with inert gas core ____________________________________________
Be, [He] 2s2
Mg, [Ne] 3s2
Ca, [Ar] 4s2
Sr, [Kr] 4s2
Ba, [Xe] 6s2
Ra, [Rn] 7s2
___________________________________________________
Because of this similarity in electronic configuration, they closely resemble in their physical and chemical properties.
1. Atomic volume, atomic and ionic radii:
Because of the addition of an extra shell of electrons to each element from Be to Ra, the atomic volume increases from Be to Ra.
The atoms of these elements although fairly large, are smaller than those of the corresponding elements of the group–1. This is due to higher nuclear charge of these atoms which tends to draw the orbital electrons towards the nucleus. The ions are also large but smaller than those of the elements of group–1. This is again due to the fact that removal of the two valence electrons to from M2+ ions increases the effective nuclear charge which pulls the electrons inwards and thus reduces the size of the ion. Atomic as well as ionic radii increases on moving down the group on account of presence of an extra shell at each step.
Elements Be, Mg Ca Sr Ba Ra
Atomic volume (c.c.) 4.90 13.97 25.9 35.44 36.7 8.00
Atomic radii (A.) 1.12 1.60 1.97 2.15 2.22 –
Radii of M2+ ions (A.) 0.31 0.65 0.99 1.13 1.35 1.40
2. Ionization energy:
Since the alkaline earth metals possess smaller size and greater nuclear charge than the alkali metals, the electrons are more tightly held and hence the first ionisation energy is greater than that of the alkali metals.
The second ionisation energy is much higher (nearly double) than the first ionisation energy.
Elements Be Mg Ca Sr Ba
IE1(eV) 9.3 7.6 6.1 5.7 5.2
IE2(eV) 18.2 15.0 11.8 11.0 10.1
Electropositive character increasing
(A) The first ionisation energy represents the energy required to remove an electron from a neutral atom (M) while the second ionisation energy represents the energy required to remove an electron from positive ion (M+) which , of course, is difficult than the former case.
(B) After the removal of one electron, the effective nuclear charge increases and hence the remaining electron is held even more tightly leading to very high ionisation energy.
(C) However the second ionization energy of the second group elements is less than the second ionization energy of the 1st group element.
On the basis of the fact that the M2+ ions are extensively hydrated and a large amount of energy, known as hydration energy is released in the process which counterbalances the higher value of IE2 which is responsible for the formation of M2+ ions in solutions. Higher value of hydration energy, in turn, is due to their much larger charge to size ratio leading to a much stronger electrostatic attraction on the oxygen of water molecule surrounding them. Formation of M2+ ions of alkaline earth metals produce stable inert gas configuration.
3. Oxidation state:
Due to the presence of two s-electrons in the outermost orbit, high heat of hydration of the dipositive ion and comparatively low values of IE2, the alkaline earth metals are bivalent. Moreover, since the bivalent ions have an inert gas configuration, it is very difficult to remove the third electron from the element and hence oxidation state higher than 2 are not encountered. Further, the bivalent ion has no unpaired electron hence their compounds are diamagnetic and colourless provided their anions are also colourless. Thus in short, it can be said that the chemistry of these elements is the chemistry of dipositive ions.
4. Nature of bonds:
Like alkali metals, all alkaline earth metals (except Be) form predominantly ionic compounds. Most of these ionic compounds are soluble in water. Beryllium forms covalent compounds due to their smaller atomic and ionic radii, hence beryllium compounds are soluble in organic solvents.
5. Electropositive character:
Due to their large size and comparatively low ionisation energies, the alkaline earth metals are strongly electropositive elements. However, these are not as strongly electropositive as the alkali metals because they have smaller size and higher ionization energies than the alkali metals, and hence, unlike alkali metals, these elements do not emit electrons on exposure to light. On moving down the group, the electropositive character increases.
6. Electronegativity.
Since these elements are electropositive, they have low electronegativities which decrease on moving down the group. The relatively high electronegativity value of Be is due to its small size.
7. Metallic character. The alkaline earth elements are metals and their metallic character increases down the group.
Metallic nature is again due to low ionization energy and since the ionization energy decreases down the group, their metallic character increases. Further due to larger number of valence electrons and HCP or CCP type of metallic packing, they are less soft, less malleable and more hard than their corresponding alkali metal.
This packing leads to stronger cohesive forces and hence these are denser than the alkali metals because they can be packed more tightly due to their greater charge and smaller radii. However, the difference in crystal structure and variations in the rate of change of atomic weight as compared to atomic radii.
This packing leads to high melting points than those of alkali metals. This is because these metals possess two valence electrons which are more strongly bonded in the solid state than the alkali metals.
8. Heat and electrical conductivity.
Due to the presence of two loosely held valence electrons (per atom) which are free to move throughout the crystal structure, the alkaline earth metals, are good conductors of heat and electricity.
9. Flame coloration.
Salts of alkaline earth metals, except that of Be and Mg, produce characteristic colour to the flame due to easy excitation of electrons to higher energy levels. Beryllium and magnesium atoms due to their small size, bind their electrons more strongly i.e. their ionisation energies are high. Hence they possess high excitation energy and are not excited by the energy of the flame to higher energy state with the result no colour is produced to the flame.
Ca = Brick orange Red Sr = Red Ba = Apple Green
10. Reducing properties (oxidation potential).
Alkaline earth metals have two electrons outside the noble gas configuration. Due to their large size and low ionisation energy, they can easily lose outermost two electrons and hence undergo oxidation (loss) of electrons easily.
M → M2+ + 2e–
They are, therefore strong reducing agents. Further since the ionization energy decreases (i.e. electropositive character increases) from Be to Ba, the reducing character of the alkaline earth elements increases from Be to Ba, which is evident from the value of their oxidation potential which increases from Be to Ba (recall that higher the value of oxidation potential or higher the negative value of reduction potential, greater is the electropositive character and reducing property of the element).
Element Oxidation reaction Oxidation potential
Be Be → Be2+ + 2e– 1.69 volt
Mg Mg → Mg2+ + 2e– 2.37 volt
Ca Ca → Ca2+ + 2e– 2.87 volt
Sr Sr → Sr2+ + 2e– 2.89 volt
Ba Ba → Ba2+ + 2e– 2.90 volt
Be, being very small in size has a very large I.P. which is not counter balanced by the hydration energy. Hence beryllium, on account of its relatively lower oxidation potential, liberates hydrogen from acids slowly; on the other hand, other elements having high values of oxidation potential react vigorously even with water.
