S-BLOCK ELEMENTS OF THE MODERN PERIODIC TABLE. What are S-BLOCK ELEMENTS ? Their POSITION in the Periodic Table and their PROPERTIES..!

S-BLOCK ELEMENTS OF THE MODERN PERIODIC TABLE

S-block Elements

The s-block of the periodic table consists of the first two groups i.e the groups IA and IIA namely the Alkali Metals and the Alkaline Earth Metals alongwith Hydrogen and Helium. 

The Group IA of the Periodic Table consists of elements like Lithium (Li), Sodium (Na), Potassium(K), Rubidium (Rb), Cesium (Cs) and Francium (Fr) which are collectively known as Alkali Metals.These are so called because they form hydroxides on reaction with water which are strongly alkaline in nature.

The Group IIA consists of elements like Beryllium(Be),Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium(Ra) which are collectively known as Alkaline Earth Metals (except Beryllium) .They are so called because their oxides and hydroxides are

alkaline in nature and these metal oxides are found in the earth's crust. 

In the s-block elements the last electron enters the outermost s-orbital and as the s-orbital can accommodate only two electrons, that is why only two groups i.e group 1 & 2 belong to the s-block of the Periodic Table. Thus the outermost orbital of s block elements consists of one or two electrons and the orbital next to the outermost shell i.e the penultimate shell has either 2 or 8 electrons. This is the reason why the s-block elements show a fixed valency which depends on the number of electrons present in the outermost shell. The last element of both the groups i.e, Francium(Fr) and Radium(Ra) are radioactive and possess different properties than the other elements of the same group.

Characteristic Properties of elements in s-blocks of Modern Periodic Table :-

1) Electronic Configuration :- The general electronic configuration of s-block elements is ns 1 for Alkali metals and ns 2 for Alkaline earth metals where n = 2 to 7 . All the Alkali metals have one valence electron and these loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. The Alkaline Earth Metals have two electrons in the s-orbital of the valence shell. Like Alkali Metals, these elements are also very electropositive.

2) Metallic character :- All the Alkali metals are silvery white, soft and light metals. The Alkaline Earth metals, in general,are silvery white, lustrous and relatively soft but harder than the Alkali metals.Beryllium and magnesium appear to be somewhat greyish. The Metallic character increases as we go down both groups. Both the Alkali metals and the Alkaline Earth metals are highly malleable and ductile and have a very high tendency to loose electrons to formpositive ions and hence they are highly electropositive.

3) Atomic Density :- The Alkali metals and the Alkaline Earth metals both have low density. This is because they have large ionic size due to which their atomic nuclei are widely separated in their crystal lattices. The density increases down both groups.

4) Melting and Boiling Points :- The melting and boiling points of the Alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them. The melting and boiling points of Alkaline Earth metals are higher than the corresponding alkali metals due to their smaller sizes. The trend is, however, not systematic but it slightly decreases down the group.

5) Oxidation state :- The Alkali metals show only +1 oxidation state, while Alkaline Earth metals show +2 oxidation state only.

6) Atomic and Ionic radii :- Both the Alkali metals and the Alkaline Earth metals have large atomic and ionic radii. The Atomic and the Ionic radii increases as we move down both groups. But as we go from group I to group II in the same period the Atomic and the Ionic radii decreases.

7) Electrode potential :- The Alkali metals are strong reducing agents. The Standard Electrode Potentials of all alkali metals lie between -2.7V and -3.0V, indicating a strong tendency to form cations in solution. The Alkaline Earth Metals also have negative values of their Standard Electrode potentials.

8) Ionisation Energies :- The ionization energies of the Alkali metals are considerably low and decreases down the group from Lithium (Li) to Cesium (Cs).This is due to the increasing size, increasing nuclear charge, and the outermost electron is very well screened from the nuclear charge. The Alkaline Earth metals have low ionization energies due to their large size of the atoms. Since the atomic size increases down the group, their ionization energies decreases down the group. But as we go from group I to group II in the same period Ionisation Energies increases. The first ionisation energies ( means the energy required to remove the first electron from the atom ) of the alkaline earth metals are higher than those of the corresponding Group I metals. This is due to their small size as compared to the corresponding alkali metals. But the second ionisation energies (means the energy required to remove the second electron from the atom) of the alkaline earth metals are smaller than those of thecorresponding alkali metals.

9) Magnetic Properties :- Alkali metals are attracted by the applied magnetic field and hence are Paramagnetic in nature whereas the  Alkaline Earth metals are repelled by the magnetic field and hence are diamagnetic in nature.

10) Complex Formation :- Both the Alkali metals and the Alkaline earth metals show weak tendency to form complexes because they have no low energy vacant orbital available for bonding with lone pair of ligands.This is due to large size, low nuclear charge and poor ability to attract electrons.

11) Chemical Reactivity :- The alkali metals are highly reactive metals and the reactivity increases down the group. The reactivity is due to :-

(a) Low value of first ionization energy

(b) Large size

(c) Low heat of atomization 

Reactivity :- 

Reactivity of Alkali Metals :-

1) Alkali metals lose their lustre very easily upon exposure to air.

2) Alkali metals react with water and form hydroxides along with liberation of hydrogen.

3) Alkali metals react with hydrogen and form ionic hydrides and these metal hydrides in turn react with water to give back the hydrogen.

4) Alkali metals react with halogens very easily forming halides.

5) All alkali metals are readily soluble in liquid ammonia.

6) The alkali metals are good reducing agents and the reducing property increases from Li to Cs.

7) Alkali metals can form alloys with other elements in the same period or with metals in other groups.

8) Alkali metals react with sulphur and phosphorus on heating to form sulphides and phosphides.

9) Alkali metals combine with mercury to form amalgams.

Reactivity of Alkaline Earth Metals :-

The alkaline earth metals are less reactive than the alkali metals. 

Reactivity of alkaline earth metals increases as we move down the group.

1) Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.

2) All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.

3) All the elements except beryllium combine with hydrogen upon heating to form their hydrides.

4) The alkaline earth metals readily react with acids liberating dihydrogen.

5) Like alkali metals, the alkaline earth metals are strong reducing agents.

6) Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.

7) The alkaline earth metals burn in oxygen to form the monoxide.

8) The alkaline earth metals also form salts of oxoacids like carbonates,sulphates and nitrates.

9) They directly combine with carbon and form carbides .

12) Conductivity :- Both the Alkali metals and the Alkaline Earth metals are good conductors of heat and electricity.

13) Colour :- All the alkali elements are silvery white solid. When freshly cut, they have a bright lusture which quickly fades due to surface oxidation.The silvery luster of alkali metals is due to the presence of highly mobile electrons of the metallic lattice.

The alkaline earth metals, in general,are silvery white and lustrous .

14) Flame colouration :- This property is due to the ease of excitation of the valence electrons. When elements or their compounds are introduced to flame, the electrons absorbs energy from the flame and gets excited to higher energy levels. When these electrons return to their ground state, they emit absorbed energy in form of visible light having characteristic wavelengths. Depending upon the wavelength of light emitted, different colours are imparted to the flame. Salts (generally chlorides) impart characteristic colours to the Bunsen flame. 

The alkali metals and their salts impart a characteristic colour to flame. 

For eg :-

Lithium - Crimson Red

Sodium - Golden Yellow

Potassium - Pale Violet

Rubidium - Violet

The alkaline earth metals also imparts colours to flame. 

For eg :-

Strontium - Bright crimson

Calcium - Brick red

Barium - Apple green

Radium - Carmine – red

15) Photoelectric Effect :- Alkali metals (except Li) and the Alkaline Earth metals exhibit photoelectric effect (A phenomenon of emission of electrons from the surface of metal when light falls on them). The ability to exhibit photoelectric effect is due to low value of ionization energy of alkali and alkaline earth metals. Li does not emit photoelectrons due to high value of ionization energy.

Description of Helium (group I element) :-

Helium was Discovered by Sir Ramsey in 1895. Helium is an element with symbol He and atomic number 2. It is a colorless, odourless, tasteless, non-toxic, inert and monatomic gas. It is the second lightest element and is the most abundant element in the universe.

Helium is the least water soluble monoatomic gas. It's the less reactive element and doesn't essentially form chemical compounds. The density and viscosity of helium vapour are very low.It has low boiling point, low density and high thermal conductivity.Natural gases contain

higher helium concentrations than the atmosphere.

Description of Hydrogen (group II element) :-

Hydrogen is an element with symbol H and atomic number 1. It is a colourless, odorless, tasteless, flammable and nontoxic gas atmospheric temperatures and pressures.It is the most abundant element in the universe. Hydrogen is the lightest of all gases.Hydrogen burns in air with a pale blue colour, almost invisible flame. Hydrogen has the highest combustion energy,This property makes it the fuel of choice for multi-stage rockets. Hydrogen has the lowest boiling point of any element except helium. Hydrogen was Discovered by H. Cavendish and named by A. Lavoisier .It's Density is about 0.08988 g/l . It's Melting point is about 14.01 K and Boiling point is about 20.28 K . It has dimagnetic magnetic properties.Hydrogen gas can form explosive mixtures with air. The destruction of the Hindenburg airship was an famous example of hydrogen combustion. Hydrogen is the only element that has different names for its isotopes i.e Deuterium , Protium and tritium . The electrolysis of water is a simple method of producing hydrogen. Hydrogen can be prepared in several different ways, but economically the most important processes involve removal of hydrogen from hydrocarbons.

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