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# S Block Elements

Alkali and alkaline Earth metals are the most violently active of all the metals. These occur in the combined form with halide, sulphate, carbonate, silicate ions, etc. and are not found in the free state in nature, as these are readily oxidized.

A number of alkali and alkaline Earth metals are found in abundance in the Earth's crust. Among these, calcium is the fifth most abundant element in the Earth's crust and hence the third most abundant metal after aluminium and iron. Vast sedimentary deposits of CaCO3 occur over large parts of the Earth's surface.

In the sixth and seventh place, the most abundant elements in crystal rocks are magnesium and sodium. Potassium (eighth place) is the next most abundant element after sodium. There is an unlimited supply of NaCl in natural brine's and oceanic waters. The occurrence of other metals in the Earth's crust is very poor. Francium and radium are radioactive and have low terrestrial abundance.

## Alkali and Alkaline Earth Metals

1. S – Block elements comprises of the first two groups, alkali metals and alkaline earth metals.
2. Group I elements are called as Alkali metals and Group ii elements are alkaline earth metals.
3. Group IA of the periodic table consists of Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs) and Francium (Fr).
4. The group – II elements are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Barium (Ba) and Radium (Ra).
5. The last element of both group, Fr and Ra are radioactive and possess different properties than the other elements of the same group.
6. All these elements are metallic in nature and are highly reactive.
7. They are collectively known as “alkali metals” as hydroxides of these metals are soluble in water and the solutions are alkaline in nature.
8. A water soluble base is known as alkali.

## Electronic Configuration

The alkali metals have one electron in the ‘s’ sub-shell of their valence shell, while alkaline earth metals have two electrons entering in the ‘s’ orbital of the valence shell. The general electronic configuration of these elements may be represented as:

[Noble gas] ns1; where n = 2 to 7.
[Noble gas] ns2; where n = 2 to 7.

Electronic configurations of alkali and alkaline earth metals

 Element Symbol Atomic number Electronic configuration Lithium Li 3 [He]2s1 Sodium Na 11 [Ne]3s1 Potassium K 19 [Ar]4s1 Rubidium Rb 37 [Kr]5s1 Cesium Cs 55 [Xe]6s1 Beryllium Be 4 [He]2s2 Magnesium Mg 12 [Ne]3s2 Calcium Ca 20 [Ar]4s2 Strontium Sr 38 [Kr]5s2

## S Block Elements Properties

The anomalous properties of S-block elements and the reasons for such behavior. Similar physical and chemical properties are shown by elements in a group.

However, the first member of each group differs from its succeeding members (called congeners). This mainly occurs because of,
• The small size of the atom and its ion.
• The high ionization energy and electronegativity.
• The non-availability of 'd' orbitals.

### Physical Properties of S block Elements

1. Physical State

All elements are silvery or greyish white, soft and light metals. S-Block elements are metals because they have low ionization energies and have a few valence shell electrons as compared to available vacant orbitals. They are highly malleable and ductile. When freshly cut, they have a bright luster which quickly tarnishes on exposure to air due to oxidation.

2. Atomic Size

The values of atomic and ionic size reveal that alkali metals and alkaline earth metals have large atomic and ionic radii. In their corresponding periods, alkali metals are the largest, followed by alkaline earth metals. Atomic radii - Li = 1.55Å, Be = 1.12 Å It can be seen that Lithium and beryllium, the first elements of both the groups have large atomic radii. Atomic radii increases further, as they move down a group.

Example- Atomic radii Potassium = 2.48 Å, while that of Calcium in the second group is 1.97Å.

3. Density

The alkali metals have low density. This is because they have large ionic size and, therefore, their atomic nuclei are widely separated in their crystal lattices. Density gradually increases in moving down from lithium towards cesium.

4. Ionization Energy

The first ionization energies of alkali metals are quite low as compared with the elements of the other groups belonging to the same period. This is because the s-electron in the valence shell being far away from the nucleus is weakly held by the nucleus. The distance of s-electron from the nucleus increases as we go down from lithium to cesium.

As a result the ionization energy decreases in going from lithium to cesium. The second ionization energy, the energy to remove a second electron from the atom is extremely high, because the second electron has to be removed from a smaller positive ion and from a filled orbital too.

Ionization energies: Li: 5.4 Na: 5.1 K: 4.3 Rb: 4.2 Cs: 3.9

5. Electropositive Character

Alkali metals are the most electropositive elements in the periodic table. As alkali metals have low ionization energies, they have a great tendency to lose electrons forming uni-positive ions.
M M+ + e-

These elements are, therefore, said to have strong electropositive character or metallic character. The electropositive character increases as we go down from Lithium to Cesium and Beryllium to Magnesium, as the size increases as we move down a group, the ionization energy decreases.

6. Oxidation State

The alkali metal atoms show only +1 oxidation state, while alkaline earth metals show +2 oxidation state only. Because of their low ionization energies, they easily lose the outermost s electron to form the uni positive ions. Once they lose the first electron, in case of alkali metals, they achieve the noble gas configuration. The same is true with respect to alkaline earth metals, when they lose the 2 electrons present in the valence shell.

7. Flame coloration

The alkali metals and their salts, when introduced into the flame, give characteristic color to the flame.

 Li Na K Rb Cs Crimson red Golden yellow Pale violet(Lilac) Red-Violet Blue

This property of the alkali metals offers a very sensitive and reliable test for alkali metals.

### Chemical Properties of S block elements

Due to their low ionization energies, alkali metals are highly electropositive and chemically reactive. Further since the ionization energies decrease with increase in atomic number, their reactivity also increase from Lithium to Cesium. The S block chemistry is nothing but the chemical properties of all S block elements.

### 1. Action with air/Oxygen

Alkali metals tarnish air due to the formation of oxide at their surface and hence they are stored in kerosene or paraffin oil.

When burnt in oxygen, Lithium forms Lithium Oxide, Li2O, Sodium forms Sodium peroxide, Na2O and other elements form superoxides also, as follows

M + O2
MO2
M = K, Rb and Cs

Alkaline earth metals form normal oxides of type MO obtained by heating the metal in the presence of oxygen or by heating their carbonates.

Ca + O2
CaO

The second group metals do not form super oxides like the alkali metals. Peroxides can be obtained though by heating the normal oxides with oxygen at high temperature.
2BaO + O2 2BaO2

### 2. Formation of Hydroxides

Both alkali metals and alkaline earth metals form hydroxides. The hydroxides of the first group metals are strongly basic in nature. Beryllium does not react with water even at elevated temperatures. Other alkaline earth metals form hydroxides by reaction of their oxides with water.

MO + H2O M(OH)2

Alkali metals react vigorously with water to form hydroxides.

## Diagonal Relationship

Some elements of second period show similarities with elements of the third period present diagonally to each other even though they belong to different groups. The similarity in properties of elements present diagonally is called diagonal relationship. This is shown below.

For example, lithium shows resemblance with magnesium, the element of group 2 Magnesium.

### Cause of Diagonal Relationship

The similarity of properties such as electronegativity, ionization energy, size etc. between the diagonal elements is the cause of diagonal relationship.

For example, on moving from left to right across a period, there is an increase in electronegativity, while moving down in the group there is a decrease in electronegativity. Therefore, on moving diagonally, the two opposing tendencies almost cancel out and the electronegativity values remain almost same. Thus, diagonal pairs have many similar properties.

Three important diagonal pairs, which exhibit diagonal similarities, are
Lithium - magnesium
Beryllium - Aluminum
Boron - Silicon

### Diagonal Relationship between Lithium and Magnesium

• Both Li and Mg are quite hard.
• Both LiOH and Mg(OH)2 are weak bases.
• Lithium reacts with N2 to form lithium nitride, while magnesium also reacts in a similar way. Other members of the group do not display this characteristic
$6Li + N_{2} \xrightarrow[]{Heat} 2Li_{3}N$ (Lithium nitride)
$3Mg + N_{2} \xrightarrow[]{Heat} Mg_{3}N_{3}$ (Magnesium nitride)
• Both lithium and magnesium combine with oxygen to form monoxides. Other members form peroxides and super oxides.
$4Li + O_{2} \rightarrow 2Li_{2}O$ (Lithium monoxide)
$2Mg + O_{2} \xrightarrow[]{Heat} 2MgO$
• On strong heating, the hydroxides of both metals decompose to form respective oxides
$2LiOH \rightarrow Li_{2}O + H_{2}O$
$Mg(OH)_{2} \xrightarrow[]{Heat} MgO +H_{2}O$

## Properties of Lithium

For example, lithium displays an anomalous behavior when compared to sodium and rest of the family members of the alkali metal family. The difference between the first members from its succeeding members is highlighted with lithium as a representative example,
• The first member of each group has the smallest size of atom and its ion in its group. The size goes on increasing as we go down a group.
• The first member has largest ionization energy because of small atomic size: the ionization energy decreases down the group.
• All the elements of second period have abnormally low electron affinity when compared to the second member because of its small atomic size. The electron affinity decreases as we move down a group.
• The bond strengths of the compounds depend on the size of the atom. Small size of atom results in relatively high cohesive properties associated with relatively strong inter-metallic bonding; large atoms usually form weak bonds. Thus bond strengths of the compounds decrease as we move down the group.
• For example, lithium has relatively high heat of atomization, melting and boiling points, density and hardness.
• The first member will have a higher electronegativity than other group members and shows greater property to form covalent bonds.
• Lithium halides are covalent while halides of other members of group 1 are ionic in nature.
• Due to very small size of the cation of second period (e.g. Li+, Be2+) the first member has high hydration energy.
• For example, Li+ has very high hydration energy and as a result, acts as an excellent reducing agent in aqueous solution.
• There are no vacant 'd' orbitals in the valence shell of the first member of the group. It is thus limited to a four co-ordination number i.e., only four electron pairs. On the other hand, the elements of third period (second member of the group) have vacant '3d' orbitals in its valence shell.
• Therefore, these can have more than four co-ordination number and can extend their octets. For example, beryllium forms BeF42- while aluminium forms AlFS63- in solution.