What are S-Block Elements?

Elements for which the last electron or valence electrons enter the s-subshell are known as s-block elements. They are electropositive in nature and lose electrons readily. The s-block contains the elements of Groups 1 and 2 in the periodic table. Group 1 contains alkali metals and Group 2 contains alkaline earth metals. Highly reactive in nature, they do not exist in the free state. They are found in the form of ionic compounds. Group 1 elements have only one electron in the valence shell while Group 2 elements have two electrons in their valence shell. They tend to lose a single electron and achieve a noble gas configuration. Therefore, they are highly electropositive in nature.

Alkali Metals in Periodic Table   

In the periodic table, alkali metals belong to Group 1. The electronic configuration of all alkali metals in the periodic table is shown below.

All alkali metals contain one valence electron. Due to this property, s-block elements of Group 1 show comparable properties.

Properties of Alkali Metals

The properties of alkali metals depend upon the size and binding of s electrons. Based on these two factors, the properties of alkali metals can vary in the following ways:

Size

The size of alkali metals increases while moving down the group because the number of shells increases. Due to this atomic radius down the group increases. In Group 1, lithium (Li) is the smallest and cesium (Cs) is the largest in size.   

Ionization Energy

Ionization energy is the amount of energy required to remove the electron from an isolated atom. Alkali metals have only one electron in the s orbital. They can either lose an electron or gain one electron to attain the noble gas configuration. So, their first ionization energy is low. They are electropositive in nature and form positively charged ions. After the removal of one electron, their oxidation state changes from 0 to +1. Group 1 elements show a +1-oxidation state.  

${\mathrm{M}}_{\left(\mathrm{g}\right)}\to {{\mathrm{M}}_{\left(\mathrm{g}\right)}}^{+} + {\mathrm{e}}^{-}$

where $\mathrm{M} = \mathrm{Na}, \mathrm{K}, \mathrm{Li}, \mathrm{Rb}, \mathrm{Cs}$

When one electron is removed from the valence shell. The ions formed are stable due to noble gas configuration. So, the second ionization energy is quite high.  

While moving down the group, the atomic size increases due to an increase in the number of shells. As the shell number increases, the distance between the nucleus and valence shell increases. The nucleus does not have a stronghold at large distance. So, the removal of electrons from such atoms is easier. Thus, ionization energy decreases down the group.   

Reducing Properties

Alkali metals act as extraordinarily strong reducing agents due to their ability to lose electrons. They have highly negative electrode potential values. Reducing nature is determined from the standard electrode potential values. They reduce the acids and water very easily and form H_2.  

Li has the highest reduction potential value among the all-alkali metals. So, it is the strongest known reducing agent.  

Coloration to the Flame

They give out characteristic color when exposed to the flame. When they are heated in the Bunsen burner with the help of platinum wire, electrons from the shells get excited and move to the higher energy states. But life in the excited state is truly short. So, they return to the ground state by emitting energy. This emitted energy is the reason behind the coloration.  

Ionic Compound Formation

As it is known that s- block elements are highly electropositive, so they readily combine with electronegative p-block elements. The atoms have significant differences in the electronegativities, so the bond formed between them is ionic. In the ionic molecule, they have a +1-oxidation state. Examples of ionic compounds are NaCl, LiCl, KCl, CsCl, etc. Ionic compounds have high melting compounds. They are highly stable. NaCl is also known as table salt.   

Chemical Properties of Alkali Metals

In the isolated state, they are highly reactive. They do not exist in the free state. They have low ionization energies and are bigger in size. All these factors contribute to the high chemical reactivity of alkali metals.  

1) Superoxide, oxides, and peroxides 

When these elements are burned in the free air, they form oxides, peroxides, and superoxide with the metal ions. Li always forms oxide Li₂O_3. Na forms the peroxide Na₂O₂, heavier elements form superoxide.  

2) Hydroxides

Hydroxide ion has the chemical formula of OH^-. It has a –1 charge and it readily combines with positively charged ions. The resulting molecule is MOH, where M can be Li, Na, K, Na, Rb, and Cs. They are strong bases. Bond is highly polar. When it is reacted with a strong acid, it forms salt and water as the products.   

3) Carbonates  

The general formula of alkali metal carbonate is M₂CO₃. Li is small in size; it cannot accommodate elements around itself. So, the carbonate of Li is highly unstable. It decomposes when heated.  

When moving down the group, the size of the elements increases, so the stability of carbonates increases.  

4) Solutions in liquid Ammonia 

When alkali metals are dissolved in the liquid ammonia, they form a blue-colored solution. This blue color is due to the free electrons. The solution has free electrons, so it is paramagnetic in nature. It is strongly reducing in nature. The Reaction between an alkali metal and liquid ammonia is shown below:  

Alkaline Earth Metals

Elements of group 2 are known as alkaline earth metals. They are called so because of their oxides. Their oxides are alkaline in nature and Earth word is given to those which remain unchanged upon heating. They do not occur in the free state. They are highly reactive. They have the general electronic configuration of ns². They have two valence electrons in the s orbital. But the noble gas shell hinders the direct attraction between the nucleus and electrons. So, electrons are held loosely. The electronic configuration is shown below.

As it is clear from the table. All alkaline earth metals contain two valence electrons. So, they have similar chemical and physical properties.  

Properties of Alkaline Earth Metals

Properties of alkali metals are discussed below. 

Size

Alkali earth metals have a smaller radius than group 1 in the periodic table. As moving down the group, the number of shells increases. Due to this atomic radius increases as moved down the group.  

Ionization Energy 

Alkaline earth metals have higher ionization energy values than alkali metals. They have 2 electrons in the s orbitals. So, the removal of electrons from a stable orbital is slightly difficult. But they can lose two electrons from the subshell and form M²⁺ ions. They attain noble gas configuration and hence are highly stable. After ionization, the oxidation state of alkaline earth metals changes from zero to +2.  

where M can be any metal.  

As moving down the group, the atomic size increases due to an increase in the number of shells. As the shell number increases, the distance between the nucleus and valence shell increases. The nucleus does not have a stronghold at larger distances. So, the removal of electrons from such atoms is easier. So, Ionization energy decreases down the group.  

Reducing Nature 

They have highly negative electrode potential values. They are comparatively smaller in size than alkali metals and have high hydration energies. So, they are better reducing agents than alkali metals.  

As moving down the group, the reduction potential values increase down the group. So, the tendency of the elements to act as a reducing agent also increases.  

Coloration to the Flame

They have a smaller size than group 1. So, magnesium and beryllium do not impart any color when brought under the flame. They have high ionization energy values. However, other elements show characteristic color in flame.  

Chemical Properties of Alkaline Earth Metals 

Elements have a smaller size and increased nuclear charge. So, they have high ionization energies. The oxides formed by alkaline earth metals have a covalent character. But hydroxides are less basic than alkali metals. Compounds of metals are insoluble in water due to their covalent character.   

Hydrides 

They form metal hydrides with the general chemical formula of MH₂. Beryllium hydride is quite a common example of metal hydride. It is quite ionic in nature and is a 3-D polymeric solid.   

Carbides

When they are heated in the furnace in the presence of carbon, they result in the formation of metal carbide with the chemical formula MC₂. They have an acetylide unit present in them, which has a triple bond. It liberates ethene upon reacting with water.  

Sulfates

Sulfates of Mg are useful in pharmaceuticals. Magnesium sulfate has the chemical formula MgSO₄.7H₂O. It is known as Epsom salt and is used as a purgative. Calcium sulfate has the chemical formula CaSO₄.2H₂O. It is known as gypsum and is heated to form plaster of Paris (POP). It is extensively used in industries.  

Formulas

Context and Applications

This topic is significant in the professional exams for both undergraduate and

graduate courses, especially for

• Bachelor of Science
• Master of Science

Practice Problems  

1) Why are s-block elements called alkali metals?  

Answer: Alkali is the other word used for the bases. They are called alkali metals because they form alkalis when they are reacted with water. Examples are NaOH and KOH. They further react with acids to form water and salts.  

2) Why are s- and p-block elements called representative?  

Answer: Representative term is used for the elements in which the valence electrons enter the s or p subshell. Their subshells are not filled completely. By losing or gaining some electrons, they attain noble gas configuration. So, they are highly reactive. This is why s and p- block is called representative elements.