CBSE Class 11 Chemistry Chapter 10: The s-Block Elements NCERT Solutions

The common physical and chemical features of alkali metals (Group 1 elements) are as follows:

Physical Properties:

1. Softness: Alkali metals are very soft and can be easily cut with a knife. For example, sodium can be readily sliced.
2. Appearance: They are typically silvery-white and have a metallic luster.
3. Low Density: Due to their large atomic sizes and weak metallic bonding, alkali metals have low densities. The density generally increases down the group, with Potassium (K) being an exception as it is less dense than Sodium (Na).
4. Low Melting and Boiling Points: The metallic bonding in alkali metals is weak, leading to low melting and boiling points compared to other metals.
5. Characteristic Flame Colors: When heated in a flame, alkali metals and their salts emit characteristic colors. This occurs because the heat excites the valence electrons to higher energy levels, and they emit energy as visible light when returning to the ground state. For example, Lithium gives a crimson red, Sodium gives a yellow, Potassium gives a violet, Rubidium gives a red-violet, and Cesium gives a blue color.
6. Photoelectric Effect: Alkali metals, particularly Cesium and Potassium, exhibit the photoelectric effect, meaning they can emit electrons when exposed to light of suitable frequency.

Chemical Properties:

Alkali metals are highly reactive due to their low ionization enthalpies, and their reactivity increases down the group.

1. Reaction with Water: They react vigorously with water to form their respective metal hydroxides and liberate hydrogen gas. The reaction becomes more spontaneous down the group. The general reaction is:

   $$2M\; +\; 2H\_2O\; \backslash longrightarrow\; 2M^+\; +\; 2OH^-\; +\; H\_2$$

2. Reaction with Dihydrogen: Alkali metals react with dihydrogen at elevated temperatures to form ionic hydrides (M⁺H^-). These hydrides are typically white crystalline solids with high melting points.

   $$2M\; +\; H\_2\; \backslash longrightarrow\; 2MH$$

3. Reaction with Halogens: Except for Lithium, all alkali metals react directly with halogens to form ionic halides (MX).

   $$2M\; +\; X\_2\; \backslash longrightarrow\; 2MX$$

   (Where M = Li, Na, K, Rb, Cs and X = F, Cl, Br, I)

   Lithium halides, however, exhibit some covalent character due to the high polarizing power of the small Li⁺ ion.

4. Reducing Agents: Alkali metals are strong reducing agents. Their reducing power increases down the group. Lithium is the strongest reducing agent among them due to its high hydration energy, which compensates for its high ionization enthalpy and low enthalpy of sublimation.

5. Solution in Liquid Ammonia: Alkali metals dissolve in liquid ammonia to form deep blue colored solutions. These solutions contain solvated metal cations ([M(NH₃)_x]⁺) and solvated electrons ([e(NH₃)_y]^-). The solvated electrons are responsible for the characteristic blue color and electrical conductivity.

   M + (x+y)NH_3 \longrightarrow [M(NH_3)_x]^+ + [e(NH_3)_y]^-

   These solutions are paramagnetic. On standing, they can slowly liberate hydrogen gas, forming amides:

   [M(NH_3)_x]^+ + e^- + NH_3 \longrightarrow MNH_2 + H_2 + xNH_3

   In concentrated solutions, the blue color may change to bronze, and the solution becomes diamagnetic.

The general characteristics and gradation in properties of alkaline earth metals (Group 2 elements) are discussed below:

General Characteristics:

1. Electronic Configuration: The general electronic configuration of alkaline earth metals is $ns^2$, where 'n' represents the principal quantum number. This indicates they have two electrons in their outermost s-orbital.
2. Oxidation State: Alkaline earth metals readily lose their two valence electrons to achieve the stable electronic configuration of the preceding noble gas. Therefore, they exhibit a characteristic oxidation state of +2 in their compounds.
3. Atomic and Ionic Radii: Alkaline earth metals have smaller atomic and ionic radii compared to alkali metals of the same period. As one moves down the group (from Be to Ra), the principal quantum number increases, leading to an increase in atomic and ionic radii. This is because the added electrons occupy new shells, and although the nuclear charge increases, the shielding effect of inner electrons also increases, resulting in a net expansion of the atomic size.

Gradation in Properties:

As we move down the group from Beryllium (Be) to Radium (Ra), the properties of alkaline earth metals show a distinct gradation:

1. Atomic and Ionic Radii: Increase down the group due to the addition of new electron shells.
2. Ionization Enthalpy: Generally decreases down the group. This is because the valence electrons become farther from the nucleus and are more effectively shielded by inner electrons, making them easier to remove. However, the ionization enthalpies of alkaline earth metals are higher than those of alkali metals in the same period because of their smaller size and higher effective nuclear charge.
3. Electronegativity: Decreases down the group.
4. Metallic Character: Increases down the group as the tendency to lose electrons increases.
5. Reactivity: Increases down the group. They are strong reducing agents, but less so than alkali metals. Their reactivity is primarily due to their tendency to lose two electrons to form +2 ions.
6. Solubility of Hydroxides and Salts: The solubility of hydroxides (like Be(OH)₂, Mg(OH)₂) and salts like sulfates (e.g., BeSO₄, MgSO₄) generally increases down the group. However, the solubility of carbonates and phosphates decreases down the group.
7. Thermal Stability of Carbonates and Bicarbonates: The thermal stability of carbonates and bicarbonates increases significantly down the group. For example, Beryllium carbonate is unstable and decomposes easily, while Barium carbonate is very stable.