The s-block Elements- Introduction, Properties, Examples

The \({\rm{s – }}\)block elements are the elements of the periodic table that belong to groups \(1\) and \(2.\) Group \(1\) elements are referred to as alkali metals, while group \(2\) elements are alkaline earth metals. Because the oxides and hydroxides are alkaline, they are given this name. The alkali metals have one \({\rm{s – }}\)electron in the valence shell of their atoms, while the alkaline earth metals have two \({\rm{s – }}\)electrons. These metals are both extremely reactive and can generate mono-positive and di-positive ions. In this article, we will explore everything about the characteristics of group \(1\) and group \(2.\)

General Characteristics of s-Block Elements

Some of the important characteristics of \({\rm{s – }}\)block elements are as follows:

1. Occurrence- The elements do not occur in a state in nature. They are always found in the combined state, usually carbonates, sulphates, silicates, phosphates, etc.
2. Abundance in nature- Among the elements of group \(1,\), Sodium and potassium are the seventh and eighth most abundant element in the earth’s crust. Other elements are less abundant.
3. Among the elements in group \(2,\), beryllium is not much abundant, but magnesium and calcium are quite abundant.
4. Electronic configuration- These elements possess electronic configuration of the type \({\rm{n}}{{\rm{s}}^{{\rm{1 – 2}}}}.\) All the inner shells in these elements are completely filled.
5. Valence electrons– The s-block elements possess either \(1\) or \(2\) electrons in their outermost shell.
6. Oxidation state– The elements of group \(1\) always exist in the \(+1\) oxidation state, while those of group \(2\) in the \(+2\) oxidation state.
7. Nature of compounds– The compounds of \({\rm{s – }}\)block elements are predominantly metallic in nature.
8. Anomalous behaviour- The first elements of both the groups show weird behaviour and differ from other elements in many ways. The behaviour is shown due to the following three factors-
9. i. Small size- The atoms and ions are much smaller than other elements of the group.
10. ii. High electronegativity- The first element is the most electronegative.
11. iii. Non-availability of \({\rm{d – }}\)orbitals- The higher members of the group belonging to higher periods possess \({\rm{d – }}\)orbitals and may utilize them for bond formation.
12. Diagonal Relationship- The first elements of groups often resemble the second element of the neighbouring group on the right, i.e., the element placed diagonally opposite to it. This type of resemblance is called a diagonal relationship. For example, lithium resembles magnesium, and beryllium resembles Aluminium, as shown-

Group 1 Elements- Alkali Metals

This group consists of six elements (excluding hydrogen). These elements are lithium \(\left( {{\rm{Li}}} \right){\rm{,}}\) sodium \(\left( {{\rm{Na}}} \right){\rm{,}}\) potassium \(\left( {{\rm{K}}} \right){\rm{,}}\) rubidium \(\left( {{\rm{Rb}}} \right){\rm{,}}\) caesium \(\left( {{\rm{Cs}}} \right){\rm{,}}\) and francium \(\left( {{\rm{Fr}}} \right){\rm{.}}\)

General Characteristics of Alkali Metals

1. Electronic Configuration: The general electronic configuration of alkali metals is [noble gas] \({\rm{n}}{{\rm{s}}^{{\rm{ – 1}}}}\) Where \({\rm{n}}\) represents the valence shell.
2. Atomic and Ionic Radii: Alkali metals atomic and ionic radii increase as they move down the group, i.e., their size increases from Li to Cs. Alkali metals lose one valence electron to generate monovalent cations. As a result, the cationic radius is smaller than the parent atom.
3. Ionization Enthalpy: The ionization enthalpies of alkali metals are generally low in the respective period, decreasing from Li to Cs as the group progresses. Because alkali metals have such large atomic sizes, the valence \({\rm{s – }}\)electron can be easily removed. Because of larger atomic radii and the screening effect, the amount of the force of attraction with the nucleus decreases as the group progresses.
4. Hydration Enthalpy: The smaller the ion, the greater its tendency to become hydrated, and thus the higher the hydration enthalpy. Alkali metal ions’ hydration enthalpies fall as ionic sizes grows larger. Hydration enthalpy: \( \to {\rm{L}}{{\rm{i}}^{\rm{ + }}}{\rm{ > N}}{{\rm{a}}^{\rm{ + }}}{\rm{ > }}{{\rm{K}}^{\rm{ + }}}{\rm{ > R}}{{\rm{b}}^{\rm{ + }}}{\rm{ > C}}{{\rm{s}}^{\rm{ + }}}{\rm{.}}\)

Physical Properties

1. Alkali metals are all silvery-white, delicate, and light.
2. They have a relatively low density that increases on moving along the group.
3. They show \({\rm{a + 1}}\) oxidation state.
4. The alkali metals possess low melting and boiling points. The melting and boiling points decrease on moving down the group.
5. Alkali metals (except Li) exhibit a photoelectric effect.
6. They give an oxidizing flame colour. The outermost orbital electron is excited to a higher energy level by the heat from the flame. When an excited electron returns to its ground state, radiation is emitted in the visible range.

Chemical Properties

1. Reaction with air: When alkali metals are exposed to the air surface, they get tarnished due to the formation of oxides and hydroxides.
   When alkali metals are heated, they mix with oxygen to generate a variety of oxides, depending on their nature
   \({\rm{4Li + }}{{\rm{O}}_{\rm{2}}}{\rm{ + heat}} \to {\rm{2L}}{{\rm{i}}_{\rm{2}}}{\rm{O}}\)
   \(2{\rm{Na}} + {{\rm{O}}_2} + {\rm{heat}} \to {\rm{N}}{{\rm{a}}_2}{{\rm{O}}_2}\)
   \({\rm{K}} + {{\rm{O}}_2} + {\rm{heat}} \to {\rm{K}}{{\rm{O}}_2}\)
2. Reaction with water: Alkali metals react readily and violently with water to form hydroxides with the liberation of hydrogen gas.
   \(2{\rm{Na}} + 2{{\rm{H}}_2}{\rm{O}} \to 2{\rm{NaOH}} + {{\rm{H}}_2}\)
3. Reaction with Dihydrogen: Alkali metals react with hydrogen gas at about \({\rm{673}}\,{\rm{K}}\) (Li at \({\rm{1073}}\,{\rm{K}}\)) to form saline hydrides having high melting points.
   \(2{\rm{Na}} + {{\rm{H}}_2} \to 2{\rm{NaH}}\)
4. Reaction with halogens: Alkali metals readily combine with halogens to form halides.
   \(2{\rm{Na}} + {\rm{C}}{{\rm{l}}_2} \to 2{\rm{NaCl}}\)
   The reactivity of alkali metals towards a particular halogen increase on going from Li to Cs, while reactivity of halogens towards a particular alkali metal follows the order \({{\rm{F}}_{\rm{2}}}{\rm{ > C}}{{\rm{l}}_{\rm{2}}}{\rm{ > B}}{{\rm{r}}_{\rm{2}}}{\rm{ > }}{{\rm{I}}_{\rm{2}}}\)
5. Reducing nature– Alkali metals act as strong reducing agents due to their low values of ionization energies. The reducing character of alkali metals in aqueous solution has been observed to follow the sequence \({\rm{Na < K < Rb < Cs < Li}}{\rm{.}}\) This is because Li has the lowest electrode potential.
6. Solubility in liquid ammonia: Deep blue solution is created when alkali metals dissolve in liquid ammonia. The solution is conducting in nature.
   \({\rm{M + (x + y)N}}{{\rm{H}}_{\rm{3}}} \to {\left[ {{\rm{M}}\left( {{\rm{N}}{{\rm{H}}_{\rm{3}}}} \right){\rm{x}}} \right]^{\rm{ + }}}{\rm{ + }}{\left[ {{\rm{e}}\left( {{\rm{N}}{{\rm{H}}_{\rm{3}}}} \right){\rm{y}}} \right]^{\rm{ – }}}\)
   When light strikes the ammoniated electrons, they absorb energy that corresponds to the red colour, and the light that emerges is blue. The colour of the concentrated solution changes from blue to bronze. The highly concentrated solutions are diamagnetic and colourless, while the blue solutions are paramagnetic.

Uses of Alkali Metals

Some of the uses of Alkali metals are given as follows:

Group 2 Elements- Alkaline Earth Metals

Group \(2\) of the periodic table consists of six elements-beryllium \(\left( {{\rm{Be}}} \right){\rm{,}}\) magnesium \(\left( {{\rm{Mg}}} \right){\rm{,}}\) calcium \(\left( {{\rm{Ca}}} \right){\rm{,}}\) strontium \(\left( {{\rm{Sr}}} \right){\rm{,}}\) barium \(\left( {{\rm{Ba}}} \right){\rm{,}}\) and radium \(\left( {{\rm{Ra}}} \right){\rm{.}}\)

General Characteristics of Alkaline Earth Metals

The general characteristics of alkaline earth metals are described as follows:

1. Electronic Configuration: The general electronic configuration of alkaline earth metals is [noble gas] \({\rm{n}}{{\rm{s}}^{\rm{2}}}{\rm{,}}\) where n represents the valence shell.

2. Atomic and ionic radii: Alkaline earth metals have a smaller atomic and ionic radius than alkali metals. The atomic and ionic radii within the group increase as the atomic number increases. This is because these elements only contain two valence electrons, and the nucleus-valence electron force of attraction is quite small.

3. Ionization enthalpy: Due to the relatively larger size of the atoms, these metals have low ionization enthalpies. The group ionization enthalpies are expected to decrease as atomic sizes decrease.
Alkaline earth metals have greater first ionization enthalpies than respective alkali metals due to their smaller size in comparison to alkali metals.

4. Hydration enthalpy: The hydration enthalpies of alkaline earth metal ions are higher than those of alkali metal ions due to their small size and high positive charge on ions. As a result, alkaline earth metals have a higher tendency to hydrate. Since the cationic size increases, the hydration enthalpies decrease down the group.
\({\rm{B}}{{\rm{e}}^{{\rm{2 + }}}}{\rm{ > M}}{{\rm{g}}^{{\rm{2 + }}}}{\rm{ > C}}{{\rm{a}}^{{\rm{2 + }}}}{\rm{ > S}}{{\rm{r}}^{{\rm{2 + }}}}{\rm{ > B}}{{\rm{a}}^{{\rm{2 + }}}}\)

5. Metallic Character: When compared to alkali metals from the same period, they exhibit strong metallic bonds. This is due to the alkaline earth metal’s smaller kernel size and the presence of two valence electrons in the outermost shell.

Physical Properties of Alkaline Earth Metals

1. They are harder than alkali metals.
2. Due to their small size, M.P and B.P are higher than the comparable alkali metals.
3. On moving down the group, the electropositive character increases.
4. With the exception of Be and Mg, all of these metals give the flame a distinct colour.
5. The thermal and electrical conductivity of alkaline earth metals is high.
6. Except for beryllium and magnesium, the alkaline earth metals and their salts impart characteristic colours to the flame. The colour of the flames are as follows-

Chemical Properties of Alkaline Earth Metals

1. Reaction with oxygen: Due to high reactivity, these elements are oxidized by atmospheric air and get tarnished. Beryllium and magnesium are not much affected in air. Calcium and strontium are easily tarnished, while barium burns instantaneously when exposed to air.
2. Reaction with water: These metals are less reactive to water than alkali metals because they are less electropositive.
   Magnesium reacts with steam or boiling water. The rest of the members reacted even with cold water.
   \({\rm{Mg}} + 2{{\rm{H}}_2}{\rm{O}} \to {\rm{Mg}}{({\rm{OH}})_2} + {{\rm{H}}_2}\,\,\,\,\,{\rm{Ca}} + 2{{\rm{H}}_2}{\rm{O}} \to {\rm{Ca}}{({\rm{OH}})_2} + {{\rm{H}}_2}\)
3. Reaction with halogens: The alkaline earth metals combine with halogens at higher temperatures to form halides of the type \({\rm{M}}{{\rm{X}}_{\rm{2}}}{\rm{.}}\)
   \({\rm{Be}} + {\rm{C}}{{\rm{l}}_2} \to {\rm{BeC}}{{\rm{l}}_2}\)
4. Reaction with acids: These elements react with dilute acids to form corresponding salts with the liberation of hydrogen gas.
   \({\rm{Mg}} + 2{\rm{HCl}} \to {\rm{MgC}}{{\rm{l}}_2} + {{\rm{H}}_2}{\rm{Ca}} + 2{\rm{HCl}} \to {\rm{CaC}}{{\rm{l}}_2} + {{\rm{H}}_2}\)Reaction with alkalis: Except for beryllium, no other element reacts with any alkali.
   \({\rm{Be}} + 2{\rm{NaOH}} \to {\rm{N}}{{\rm{a}}_2}{\rm{Be}}{{\rm{O}}_2} + {{\rm{H}}_2}\)
5. Reaction with hydrogen: Except beryllium, all other alkaline earth metals combine directly with hydrogen on heating to form hydrides of the type
   \({\rm{M}}{{\rm{H}}_{\rm{2}}}\)
   \({\rm{Mg}} + {{\rm{H}}_2} \to {\rm{Mg}}{{\rm{H}}_2}\)

Uses of Alkaline Earth Metals

1. Beryllium is extensively used in the form of alloys.
2. Copper-beryllium alloys are used in the preparation of high strength springs.
3. Beryllium is used in making windows of \({\rm{X – }}\)ray tubes due to its property to absorb radiations.
4. Magnesium can be used in metallic form as well as an alloy.
5. Magnesium powder is used in flashbulbs, incendiary bombs and signals.
6. The compound of magnesium is used in toothpaste and in medicines as antacids.
7. Calcium is used to remove traces of air from a vacuum because it attracts oxygen and nitrogen.
8. Calcium oxide is used in making cement.
9. Calcium hydroxide is used as a disinfectant and in making bleaching powder.
10. Calcium compound is used in chalks, chewing gums and paper.
11. Strontium and barium are used in fireworks.
12. Radium is used in making glowing stickers and luminous paint in watches and sign boards.
13. Radioactive radium\(-223\) is used to treat prostate cancer.

Summary

In this article, we studied the alkali and alkaline earth metals and why they are called s-block elements. We studied the diagonal relationship between second and third-period elements and the anomalous behaviour of the first element of the group. Now we know the physical and chemical properties of the alkali and alkaline earth metals and their uses.

FAQs

Q.1. What are s-block elements? Mention their important characteristics.
Ans: The element in which the last electron enters into s-orbital is called \({\rm{s – }}\)block element. Some of the important characteristics of \({\rm{s – }}\)block elements are as follows-
1. Occurrence- The elements do not occur in a state in nature. They are always found in the combined state, usually as carbonates, sulphates, silicates, phosphates, etc.
2. Abundance in nature- Among the elements of group \(1,\) Sodium and potassium is the seventh and eighth most abundant element in the earth’s crust. Other elements are less abundant.
Among the elements in the group \(2,\) beryllium is not much abundant, but magnesium and calcium are quite abundant.
3. Electronic configuration- These elements possess electronic configuration of the type \({\rm{n}}{{\rm{s}}^{1 – 2}}.\) All the inner shells in these elements are completely filled.
4. Valence electrons- The \({\rm{s – }}\)block elements possess either \(1\) or \(2\) electrons in its outermost shell.
5. Oxidation state- The elements of group \(1\) always exist in \(+1\) oxidation state, while those of group \(2\) in \(+2\) oxidation state.
6. Nature of compounds- The compounds of \({\rm{s – }}\)block elements are predominantly metallic in nature.

Q.2. What do you understand about diagonal relationships?
Ans: The first element of the group often shows resemblance to the second element of the neighbouring group on the right, i.e., the element placed diagonally opposite to it. This type of resemblance is called a diagonal relationship. For example, lithium resembles magnesium and beryllium resembles Aluminium.

Q.3. The atomic radii of alkaline earth metals are smaller than those of the corresponding alkali metals. Explain why?
Ans: The atomic and ionic radii of the alkaline earth metals are smaller than those of alkali metals in the same period because of the increased nuclear charge in their elements.

Q.4. Why are the alkaline earth metals called \({\rm{s – }}\)block elements?
Ans: Alkaline earth metals are called \({\rm{s – }}\)block elements because the last electron in their electronic configuration occupies the s-orbital of their valence shells.

Q.5. Why do alkali metals not occur in a state?
Ans: Alkali metals are highly reactive, therefore they occur in a combined state and do not occur in a state.

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