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

Alkali metals, belonging to Group 1 of the periodic table, exhibit several common physical and chemical characteristics:

Physical Properties:

1. Softness: They are exceptionally soft metals, easily cut with a knife. For instance, sodium can be readily sliced.
2. Appearance: Alkali metals are typically light-colored and possess a silvery-white luster.
3. Low Density: Due to their large atomic sizes and relatively weak metallic bonding, they have low densities. Density generally increases down the group from Lithium to Cesium, with Potassium (K) being an exception, having a lower density than Sodium (Na).
4. Low Melting and Boiling Points: The metallic bonding in alkali metals is weak, resulting in low melting and boiling points compared to other metals.
5. Flame Coloration: When heated in a flame, alkali metals and their salts impart a distinct color to the flame. This occurs because the heat excites the outermost electron to a higher energy level, and upon returning to the ground state, it emits energy as visible light. For example, Lithium gives a crimson red color, Sodium gives a yellow color, Potassium gives a violet color, Rubidium gives a red-violet color, and Cesium gives a blue color.
6. Photoelectric Effect: Metals like Cesium and Potassium readily lose electrons when irradiated with light, demonstrating the photoelectric effect.

Chemical Properties:

Alkali metals are highly reactive, primarily due to their low ionization enthalpies, which decrease down the group, leading to increased reactivity.

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

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

2. Formation of Hydrides: Alkali metals react with dihydrogen at elevated temperatures to form ionic hydrides. These hydrides are crystalline solids with high melting points.

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

3. Formation of Halides: Except for Lithium, all other alkali metals react directly with halogens to form ionic halides (MX).

   $$2M\; +\; Cl\_2\; \backslash longrightarrow\; 2MCl$$

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

   Lithium halides, however, exhibit some covalent character due to the small size of the Li+ ion, which can easily polarize the halide ion.

4. Strong Reducing Agents: Alkali metals are strong reducing agents, with their reducing power increasing down the group. Lithium is an exception; it is the strongest reducing agent among alkali metals due to its high hydration energy, which compensates for its higher ionization enthalpy and sublimation enthalpy.
5. Solution in Liquid Ammonia: Alkali metals dissolve in liquid ammonia to form deep blue colored solutions. These solutions are electrically conducting due to the presence of solvated electrons ([M(NH_3)_x]+) and ammoniated electrons ([e(NH_3)_y]-).

   $$M\_\{(l)\}\; +\; (x+y)NH\_\{3(l)\}\; \backslash longrightarrow\; [M(NH\_3)\_x]^+\; +\; [e(NH\_3)\_y]^-$$

   These solutions are paramagnetic. In concentrated solutions, the color changes to bronze, and the solution becomes diamagnetic.

Alkaline earth metals, comprising Group 2 elements (Be, Mg, Ca, Sr, Ba, Ra), share common characteristics and exhibit predictable trends in their properties:

General Characteristics:

1. Electronic Configuration: Their general electronic configuration is $ns^2$, indicating two electrons in the outermost s-orbital.
2. Oxidation State: Alkaline earth metals readily lose these two valence electrons to achieve the stable electronic configuration of the preceding noble gas, resulting in a characteristic oxidation state of +2 in their compounds.
3. Atomic and Ionic Radii: Compared to alkali metals in the same period, alkaline earth metals have smaller atomic and ionic radii. This is because they have a higher effective nuclear charge. As one moves down the group, the principal quantum number increases, leading to the addition of a new electron shell. This effect outweighs the increase in nuclear charge, causing the atomic and ionic radii to increase down the group.
4. Ionization Enthalpy: Alkaline earth metals have higher first ionization enthalpies than alkali metals because of the greater effective nuclear charge and the stable $ns^2$ configuration, which requires more energy to remove an electron. However, their second ionization enthalpies are significantly lower, as removing the second electron from the stable $[Noble\; Gas]\; ns^1$ configuration is relatively easy. Ionization enthalpies generally decrease down the group due to increasing atomic size and shielding effect.
5. Metallic Character: They are metals with characteristic metallic properties. Their metallic character increases down the group as atomic size increases and ionization enthalpy decreases.
6. Reactivity: Alkaline earth metals are less reactive than alkali metals due to their higher ionization enthalpies. Their reactivity increases down the group. They react with water to form hydroxides and hydrogen gas, though less vigorously than alkali metals. For example, Magnesium reacts with steam, while Calcium, Strontium, and Barium react with cold water.
7. Formation of Compounds: They form ionic compounds, especially with highly electronegative elements. Their oxides and hydroxides are basic, with basicity increasing down the group.

Gradation in Properties:

The properties of alkaline earth metals show a distinct gradation down the group:

• Atomic and Ionic Radii: Increase down the group (Be < Mg < Ca < Sr < Ba).
• Ionization Enthalpy: Generally decreases down the group (Be > Mg > Ca > Sr > Ba).
• Hydration Enthalpy: Decreases down the group (Be^2+ > Mg^2+ > Ca^2+ > Sr^2+ > Ba^2+). This trend influences solubility and stability of hydrated salts.
• Metallic Character: Increases down the group.
• Basicity of Oxides and Hydroxides: Increases down the group (BeO, Be(OH)_2 are amphoteric; others are basic).
• Reactivity towards Water: Increases down the group.