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NCERT Solutions Class 11 Chemistry Chapter 10 The S-Block Elements– Here are all the NCERT solutions for Class 11 Chemistry Chapter 10. This solution contains questions, answers, images, explanations of the complete chapter 1 titled The S-Block Elements taught in Class 11. If you are a student of Class 11 who is using NCERT Textbook to study Chemistry, then you must come across chapter 10 The S-Block Elements After you have studied the lesson, you must be looking for answers of its questions. Here you can get complete NCERT Solutions for Class 11 Chemistry Chapter 10 The S-Block Elements in one place.

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NCERT Solutions Class 11 Chemistry Chapter 10 The S-Block Elements

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Class 11
Subject Chemistry
Book Chemistry Part I
Chapter Number 10
Chapter Name

The S-Block Elements

NCERT Solutions Class 11 Chemistry chapter 10 The S-Block Elements

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Question & Answer

Q.1: What are the common physical and chemical features of alkali metals ?

Ans : Physical properties of alkali metals are as follows.
(1) They are quite soft and can be cut easily. Sodium metal can be easily cut using a knife.
(2) They are light coloured and are mostly silvery white in appearance.
(3) They have low density because of the large atomic sizes. The density increases down the group from Li to Cs. The only exception to this is K, which has lower density than Na.
(4) The metallic bonding present in alkali metals is quite weak. Therefore, they have low melting and boiling points.
(5) Alkali metals and their salts impart a characteristic colour to flames. This is because the heat from the flame excites the electron present in the outermost orbital to a high energy level. When this excited electron reverts back to the ground state, it emits excess energy as radiation that falls in the visible region.
(6) They also display photoelectric effect. When metals such as Cs and K are irradiated with light, they lose electrons.
Chemical properties of alkali metals
Alkali metals are highly reactive due to their low ionization enthalpy. As we move down the group, the reactivity increases.
(1) They react with water to form respective oxides or hydroxides. As we move down the group, the reaction becomes more and more spontaneous.
(2) They react with water to form their respective hydroxides and dihydrogens. The general reaction for the same is given as
\( 2 \mathrm{M}+2 \mathrm{H}_{2} \mathrm{O} \longrightarrow 2 \mathrm{M}^{+}+2 \mathrm{OH}^{\ominus}+\mathrm{H}_{2}\)
(3) They react with dihydrogen to form metal hydrides. These hydrides are ionic solids and have high melting points.
\( 2 \mathrm{M}+\mathrm{H}_{2} \longrightarrow 2 \mathrm{M}^{+} \mathrm{H}^{-}\)
(4) Almost all alkali metals, except Li, react directly with halogens to form ionic halides.
\( \begin{array}{l}{2 \mathrm{M}+\mathrm{Cl}_{2} \longrightarrow 2 \mathrm{MCl}} \\ {(\mathrm{M}=\mathrm{Li}, \mathrm{K}, \mathrm{Rb}, \mathrm{Cs})}\end{array}\)
Since \( \mathrm{Li}^{+}\) ion is very small in size, it can easily distort the electron cloud around the negative halide ion. Therefore, lithium halides are covalent in nature.
(5) They are strong reducing agents. The reducing power of alkali metals increases on moving down the group. However, lithium is an exception. It is the strongest reducing agent among the alkali metals. It is because of its high hydration energy.
(6) They dissolve in liquid ammonia to form deep blue coloured solutions. These solutions are conducting in nature.
\( \mathrm{M}+(x+y) \mathrm{NH}_{3} \longrightarrow\left[\mathrm{M}\left
(\mathrm{NH}_{3}\right)_{x}\right]^{+}+\left[\mathrm{M}\left(\mathrm{NH}_{3}\right)_{y}\right]\)
The ammoniated electrons cause the blue colour of the solution. These solutions are paramagnetic and if allowed to stand for some time, then they liberate hydrogen. This results in the formation of amides.
\( \mathrm{M}_{(\mathrm{cm})}^{+}+\mathrm{e}^{-}+\mathrm{NH}_{3(l)} \longrightarrow \mathrm{MNH}_{(\mathrm{cm})}+\frac{1}{2} \mathrm{H}_{2(\mathrm{g})}\)
In a highly concentrated solution, the blue colour changes to bronze and the solution becomes diamagnetic. 
Q.2: Discuss the general characteristics and gradation in properties of alkaline earth metals.

Ans : General characteristics of alkaline earth metals are as follows.
(i) The general electronic configuration of alkaline earth metals is [noble gas] \( \mathrm{ns}^{2}\).
(ii) These metals lose two electrons to acquire the nearest noble gas configuration. Therefore, their oxidation state is +2.
(iii)These metals have atomic and ionic radii smaller than that of alkali metals. Also, when moved down the group, the effective nuclear charge decreases and this causes an increase in their atomic radii and ionic radii.
(iv)Since the alkaline earth metals have large size, their ionization enthalpies are found to be fairly low. However, their first ionization enthalpies are higher than the corresponding group 1 metals.
(v) These metals are lustrous and silvery white in appearance. They are relatively less soft as compared to alkali metals.
(vi)Atoms of alkaline earth metals are smaller than that of alkali metals. Also, they have two valence electrons forming stronger metallic bonds. These two factors cause alkaline earth metals to have high melting and boiling points as compared to alkali metals.
(vii) They are highly electropositive in nature. This is due to their low ionization enthalpies. Also, the electropositive character increases on moving down the group from Be to Ba.
(viii) Ca, Sr, and Ba impart characteristic colours to flames.
Ca – Brick red
Sr – Crimson red
Ba – Apple green
In Be and Mg, the electrons are too strongly bound to be excited. Hence, these do not impart any colour to the flame.
The alkaline earth metals are less reactive than alkali metals and their reactivity increases on moving down the group. Chemical properties of alkaline earth metals are as follows.
(i) Reaction with air and water: Be and Mg are almost inert to air and water because of the formation of oxide layer on their surface.
(a) Powdered Be burns in air to form BeO and \( \mathrm{Be}_{3} \mathrm{N}_{2}\).
(b) Mg, being more electropositive, burns in air with a dazzling sparkle to form MgO and \( \mathrm{Mg}_{3} \mathrm{N}_{2}\).
(c) Ca, Sr, and Ba react readily with air to form respective oxides and nitrides.
(d) Ca, Ba, and Sr react vigorously even with cold water.
(ii) Alkaline earth metals react with halogens at high temperatures to form halides.
\( \mathrm{M}+\mathrm{X}_{2} \longrightarrow \mathrm{MX}_{2}(\mathrm{X}=\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I})\)
(iii) All the alkaline earth metals, except Be, react with hydrogen to form hydrides.
(iv) They react readily with acids to form salts and liberate hydrogen gas.
\( \mathrm{M}+2 \mathrm{HCl} \longrightarrow \mathrm{MCl}_{2}+\mathrm{H}_{2(\mathrm{x})} \uparrow\)
(v) They are strong reducing agents. However, their reducing power is less than that of alkali metals. As we move down the group, the reducing power increases.
(vi) Similar to alkali metals, the alkaline earth metals also dissolve in liquid ammonia to give deep blue coloured solutions.
\( \mathrm{M}+(x-y) \mathrm{NH}_{3} \longrightarrow\left[\mathrm{M}\left(\mathrm{NH}_{3}\right)_{x}\right]^{+2}+2\left[\mathrm{e}\left(\mathrm{NH}_{3}\right)_{y}\right]^{-}\) 
Q.3: Why are alkali metals not found in nature ?

Ans : Alkali metals include lithium, sodium, potassium, rubidium, cesium, and francium. These metals have only one electron in their valence shell, which they lose easily, owing to their low ionization energies. Therefore, alkali metals are highly reactive and are not found in nature in their elemental state. 
Q.4: Find out the oxidation state of sodium in \( \mathrm{Na}_{2} \mathrm{O}_{2}\).

Ans : Let the oxidation state of Na be x. The oxidation state of oxygen, in case of peroxides, is –1.
Therefore,
2 (x) + 2(-1) =0
2x - 2 = 0 
2x = 2
X = +1
Therefore, the oxidation state of sodium is +1. 
Q.5: Explain why is sodium less reactive than potassium.

Ans : In alkali metals, on moving down the group, the atomic size increases and the effective nuclear charge decreases. Because of these factors, the outermost electron in potassium can be lost easily as compared to sodium. Hence, potassium is more reactive than sodium.

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