Introduction

The s-block elements that have only one electron in their s-orbital are called group one or alkali metals, whereas the s-block elements that have two electrons filling their s-orbital are called group two or alkaline earth metals.

The electrons in an atom occupy various sub-orbitals of the available energy levels in the order of the increasing energy. The last or the outermost electron may be present in the s, p, d, or f subshells. In this way, elements having their last electron in the s-subshell are called s-block elements.

The following is the general electronic configuration: 

Alkali metals: ns1

Alkaline earth metals: ns2

Group 1 of s-Block Elements

Group 2 of s-Block Elements

Group 1: Alkali Metals

• These elements have the electronic configuration ns1, where n signifies the valence shell.
• Since they easily dissolve in water and form highly alkaline soluble hydroxides, they are called alkali metals.
• There is an increase in the radii of the atoms from Li to Cs (moving down the group).
• The ionisation energy decreases as we go down from Li to Cs. This is due to the greater size of the atom from which the valence shell electron can be easily removed.

Physical Properties

• Alkali metals are light, soft, and silvery-white in colour.
• They generally have a low density which increases as we go down the group.
• They provide colour to an oxidising flame. The alkaline earth metal salts  produce the following flame colours: Calcium is responsible for the brick red colour. Crimson is the colour of strontium. Apple green is due to the presence of barium.

Chemical properties

• Hydroxide and dihydrogen gas are formed in reaction with water:

2M+H2O → 2M+ + 2OH– + H2

• Alkali metals are strong reducing agents. They follow the sequence Na < K < Rb < Cs < Li.

Reason being that they can only lose one electron from their valence shells.

• The alkali metals lose their shine in dry air because of oxide formation, which further leads to the formation of hydroxide in reaction with water. 

The following are examples of s-block elements of group 1:

1. Lithium

• Lithium is the least dense.
• In reaction with water, it gives a vigorous process. 
• Rechargeable lithium batteries are used in laptops, phones, and vehicles and the non-rechargeable ones are used in pacemakers, toys, and clocks.

2. Sodium

• Sodium loses its lustre when exposed to air.
• It has a strong reaction with water.
• It is used as a heat exchanger for nuclear reactors and a reagent in chemical factories.

3. Potassium

• Potassium is soft, silver, and lustrous. It has a low melting-point and is a good conductor.
• The colour of the flame is lavender and the colour of the vapour is green.
• It plays a key role in neural synapses and chemical reactions.
• It is used to treat hyperkalemia, high blood pressure, and stroke patients.

4. Caesium

• Caesium is a soft, gold-coloured metal. 
• It reacts fiercely with water and has a quick reaction to air. 
• It is most commonly used as a drilling fluid. It can also be used in vacuum tubes, radiation monitoring equipment, devising photoelectric cells, and optic glass preparation.

Group 2: Alkaline Earth Metals

• Alkaline earth metals are found in the earth’s crust and are alkaline.
• Their general electronic configuration is ns2.
• Within the group, as the atomic number increases, there is an increase in the radii of the atoms as the attraction with the nucleus becomes small and there are only two valence electrons. Hence, their ionisation energies become low and continue to decrease in the same way.
• The hydration energies decrease as we go down the group:

Be2+ > Mg2+ > Ca2+ > Sr2+ > Ba2+

Physical Properties

• Compared to alkali metals, alkaline earth metals are harder.
• Due to their small size, their melting and boiling points are higher.
• Their tendency to readily lose electrons is good and increases as we go down the group.
• They provide a separate colour to the flame, except for Be and Mg.
• They have high electrical and thermal conductivity.

Chemical Properties

• The oxygen reactions increase as we go down the group.
• They form the halides in reaction with halogens: MX2.
• The formation of the metal hydrides takes place in reaction with hydrogen, except for Beryllium.

M+H2 → MH2

General characteristics

• The alkaline earth metals react with oxygen to form monoxides which are very stable to heat.
• All halides of alkaline earth metals are ionic.
• The solubility decreases as we go down from Be to Ba.

The following are examples of s-block elements of group 2:

1. Beryllium

• Beryllium is soft and less dense. 
• It is used in copper or nickel alloys for making electrical contacts, electrodes, and springs. It enhances their thermal as well as electrical conductivity.

2. Magnesium

• The low density and good mechanical properties of magnesium make it a good choice for manufacturing phones, tablets, and laptops. 
• It is used to remove sulphur by adding it to molten iron and steel. 
• It is also used in flares, fireworks, and sparklers as it easily ignites in the air with a bright light.

3. Calcium

• Calcium is a silvery-white element.
• It is a soft metal that loses its lustre rapidly in reaction with air and water. Calcium metal is used as a reducing agent in preparing other metals such as thorium and uranium. 
• It is also used as an alloying agent for aluminium, beryllium, copper, lead, and magnesium alloys.

4. Strontium

• Strontium is a soft, silvery metal that burns when it comes in contact with air and reacts rapidly with water. 
• It is used to produce Fe magnets and refine Zn. 
• It is also used in fireworks and flares. 
• Glow-in-the-dark paints and plastics contain illuminated strontium.

5. Barium

• Pure barium is pale-yellow and has shining and malleability properties. (Malleability is the property of metals to be drawn into thin sheets.)
• Barium is often used for plugs and in vacuum tubes to dry and remove the oxygen.
• Its compounds are used by oil and gas industries to make drilling mud.
• Paint, bricks, tiles, glass, and rubber are all made of barium compounds.

Conclusion

The s-block of the periodic table comprises Group 1 (alkali) and Group 2 (alkaline earth) metals. Since their oxides and hydroxides are alkaline, these are called s-block elements. They can be differentiated by the fact that the alkali metals have one valence electron and the alkaline metals have two valence electrons. These are highly reactive: alkali metals form monopositive ions, whereas alkaline metals form dipositive ions.Both types of metals have nearly the same type of chemistry. However, some differences arise because of reduced atomic and ionic sizes and increased cationic charges in the case of alkaline earth metals. The industrially essential compounds of calcium are calcium oxide (lime), calcium hydroxide (slaked lime), calcium sulphate (plaster of Paris), calcium carbonate (limestone), and cement. Portland cement is an important constructional material. It is manufactured by heating a pulverised mixture of limestone and clay in a rotary kiln. The clinker that is thus obtained is mixed with some gypsum (2-3%) to give a fine powder of cement.The alkali metals are silvery-white, soft, and low melting. They are highly reactive.