Q: Q3. Why are alkali metals not found in free state in nature?

Alkali metals are not found in the free state in nature because they are highly reactive and electropositive elements. This high reactivity stems from two main factors:\n\n1. **Low Ionization Enthalpy:** Alkali metals have only one electron in their outermost shell (ns¹). This electron is loosely held by the nucleus and can be easily removed. Consequently, they have very low ionization enthalpies, meaning little energy is required to form a positive ion (M⁺).\n2. **Strong Electropositive Character:** Due to their low ionization enthalpy, alkali metals have a strong tendency to lose their valence electron and form stable positive ions. This makes them highly electropositive.\n\nBecause of their strong tendency to lose electrons, alkali metals readily react with other elements (like oxygen, halogens, sulfur, etc.) present in their surroundings to form stable ionic compounds. They are never found as free metals but always exist in combined forms, such as chlorides (e.g., NaCl, KCl), oxides, silicates, and carbonates.

Q: Q5. Arrange the following metals in the increasing order of their reducing power: Li, Na, K, Rb, Cs.

The increasing order of reducing power for the given alkali metals is:\n\n**Na < K < Rb < Cs < Li**\n\n**Explanation:**\n\n1. **General Trend:** For alkali metals, the reducing power generally increases down the group from Na to Cs. This is because as we move down the group, the atomic size increases, and the ionization enthalpy decreases. A lower ionization enthalpy means the electron can be removed more easily, making the element a stronger reducing agent.\n So, the expected order based on ionization enthalpy would be Na < K < Rb < Cs.\n\n2. **Anomalous Behavior of Lithium:** Lithium is an exception to this general trend. Despite having the highest ionization enthalpy among the alkali metals, it is the strongest reducing agent in aqueous solution. This is attributed to the exceptionally high hydration enthalpy of the small Li⁺ ion. The small size of Li⁺ leads to a very high charge density, causing it to be extensively hydrated in aqueous solution. The large amount of energy released during the hydration of Li⁺ (ΔH_hyd) more than compensates for its higher ionization enthalpy (ΔH_ion) and sublimation enthalpy (ΔH_sub), resulting in the most negative standard electrode potential (E°) for Lithium.\n\n The overall process for reducing power in aqueous solution involves:\n M(s) → M(g) (Sublimation)\n M(g) → M⁺(g) + e⁻ (Ionization)\n M⁺(g) + H₂O → M⁺(aq) (Hydration)\n\n The sum of these energy changes determines the standard electrode potential (E°), which is a measure of reducing power. For Lithium, the very large negative hydration enthalpy makes its E° value the most negative, indicating the strongest reducing power.\n\nTherefore, considering both the general trend and the anomaly of Lithium, the correct increasing order of reducing power is Na < K < Rb < Cs < Li.

Q: Q6. Why are alkali metals not found in free state in nature?

Alkali metals are not found in the free state in nature because they are highly reactive. They have a strong tendency to lose their single valence electron (ns¹) to form a stable cation (M⁺). This low ionization enthalpy and high electropositivity make them readily react with various elements like oxygen, halogens, and water, forming stable compounds. Due to their high reactivity, they always exist in combined forms (compounds) in nature, such as halides, oxides, silicates, etc.

Q: Q9. Why are alkali metals strong reducing agents?

Alkali metals are strong reducing agents because they have a very high tendency to lose their single valence electron (ns¹) to form a stable unipositive ion (M⁺) with a noble gas configuration. This tendency is due to:\n\n1. **Low Ionization Enthalpy:** They have the lowest first ionization enthalpies in their respective periods. This means only a small amount of energy is required to remove their outermost electron.\n2. **Large Atomic Size:** Their large atomic size results in the valence electron being far from the nucleus and experiencing less effective nuclear charge, making it easier to remove.\n3. **High Electropositivity:** Their strong tendency to lose electrons makes them highly electropositive.\n\nBy losing an electron, alkali metals themselves get oxidized (M → M⁺ + e⁻) and in turn, reduce other species. The ease with which they get oxidized makes them powerful reducing agents. Lithium, despite having a higher ionization enthalpy than other alkali metals, is the strongest reducing agent in aqueous solution due to its exceptionally high hydration enthalpy of the small Li⁺ ion.

Q: Q10. When an alkali metal dissolves in liquid ammonia the solution acquires different colours. Explain this type of behaviour.

When an alkali metal (M) dissolves in liquid ammonia, it forms a deep blue solution. This unique behaviour is due to the formation of ammoniated electrons and ammoniated metal ions.\n\n**1. Dissolution Process:**\nUpon dissolution, the alkali metal atom readily loses its valence electron to form a metal cation (M⁺) and an electron (e⁻). Both the cation and the electron get solvated (ammoniated) by ammonia molecules.\n\n**Equation:**\nM(s) + (x + y) NH₃(l) → [M(NH₃)ₓ]⁺ (ammoniated cation) + [e(NH₃)y]⁻ (ammoniated electron)\n\n**2. Deep Blue Colour:**\nThe deep blue colour of the solution is primarily due to the presence of the **ammoniated electrons** ([e(NH₃)y]⁻). These electrons absorb energy in the visible region of the electromagnetic spectrum, getting excited to higher energy levels. When they return to their ground state, they emit light corresponding to the blue colour, or rather, absorb light in the red/orange region, transmitting blue light. The solutions are also paramagnetic due to the presence of unpaired ammoniated electrons.\n\n**3. Conductivity:**\nThe solutions are highly electrically conductive, similar to metals, due to the presence of both ammoniated cations and ammoniated electrons.\n\n**4. Colour Change at Higher Concentrations:**\nAt higher concentrations (above 3 M), the blue solutions gradually change to a bronze or copper-like colour and become diamagnetic. This is because the ammoniated electrons start to associate or pair up, forming electron clusters. The metallic lustre is due to the formation of these electron clusters, which reflect light like a metal. The diamagnetic nature arises from the pairing of electrons.\n\n**5. Stability and Decomposition:**\nThe solutions are metastable and slowly decompose to form metal amides and hydrogen gas, especially in the presence of impurities or catalysts:\n2M(s) + 2NH₃(l) → 2MNH₂(ammoniated amide) + H₂(g)\nThis decomposition is faster for lithium.

Q: Q11. Why are alkali metals not found in free state in nature?

Alkali metals are highly reactive elements due to their electronic configuration and low ionisation enthalpies. They have only one electron in their outermost shell, which they readily lose to form a stable unipositive ion (M⁺). This tendency to lose an electron makes them strong reducing agents and highly electropositive. Consequently, they react vigorously with various elements like oxygen, halogens, and water present in the environment to form stable ionic compounds. Because of their high reactivity and strong tendency to form compounds, they are never found in the free or elemental state in nature; instead, they exist only in combined forms as salts, oxides, silicates, etc.

Q: Q15. Explain why alkali metals have low melting and boiling points.

Alkali metals have low melting and boiling points primarily due to the nature of their metallic bonding. Here's a detailed explanation:\n\n1. **Weak Metallic Bonding:** Alkali metals have only one valence electron per atom (ns¹). These single valence electrons are delocalised throughout the metallic lattice, forming a 'sea of electrons' that binds the positively charged metal ions (M⁺) together. However, because there is only one electron per atom contributing to this metallic bond, the overall metallic bonding forces are relatively weak.\n\n2. **Large Atomic Size:** Alkali metals have the largest atomic radii in their respective periods. The larger the atomic size, the greater the distance between the nucleus and the delocalised valence electrons. This increased distance results in weaker electrostatic attraction between the positive metal ions and the electron sea.\n\n3. **Low Cohesive Energy:** The combination of weak metallic bonding and large atomic size leads to low cohesive energy within the metallic lattice. Cohesive energy is the energy required to separate the atoms in a solid.\n\nTo melt or boil a metal, sufficient energy must be supplied to overcome these metallic bonds. Since the metallic bonds in alkali metals are weak, only a small amount of thermal energy is required to break them, resulting in low melting and boiling points. The melting and boiling points generally decrease down the group (from Li to Cs) as the atomic size increases further, leading to even weaker metallic bonding.

Q: Q16. Why are lithium salts commonly hydrated and those of other alkali ions usually anhydrous?

Lithium salts are commonly hydrated due to the exceptionally small size and high charge density of the Li+ ion. This leads to a very high charge-to-radius ratio and consequently, a very high hydration enthalpy. The strong electrostatic attraction between the small Li+ ion and the polar water molecules results in the formation of stable hydrated salts, such as LiCl·2H2O.\n\nIn contrast, other alkali metal ions (Na+, K+, Rb+, Cs+) are much larger in size and have significantly lower charge densities. Their hydration enthalpies are much smaller and are often not sufficient to overcome the lattice energy to form stable hydrates. Therefore, their salts are usually anhydrous.

Q: Q17. Why is LiF almost insoluble in water whereas LiCl is soluble not only in water but also in acetone?

The solubility of an ionic compound in water depends on the balance between its lattice energy and hydration energy. For a compound to be soluble, the hydration energy must be greater than or comparable to the lattice energy.\n\n**LiF is almost insoluble in water:**\nLiF has a very high lattice energy. Both Li+ and F- ions are exceptionally small. The strong electrostatic forces of attraction between these small, oppositely charged ions result in a very compact crystal lattice with high lattice energy. Although Li+ has a high hydration energy, the lattice energy of LiF is so high that it cannot be overcome by the hydration energy, making LiF sparingly soluble or almost insoluble in water.\n\n**LiCl is soluble in water and acetone:**\n1. **Solubility in water:** LiCl has a lower lattice energy compared to LiF because the Cl- ion is much larger than the F- ion. While the hydration energy of Li+ is still high, it is sufficient to overcome the relatively lower lattice energy of LiCl, making it soluble in water.\n2. **Solubility in acetone:** LiCl exhibits significant covalent character. According to Fajan's rules, the small size and high charge density of Li+ ion can polarize the larger electron cloud of the Cl- ion to a greater extent, leading to increased covalent character in the Li-Cl bond. Acetone is a polar organic solvent. The covalent character of LiCl allows it to dissolve in organic solvents like acetone (which is also polar) based on the principle of 'like dissolves like'.

Q: Q19. What happens when\n(a) Sodium metal is dropped in water?\n(b) Sodium peroxide dissolves in water?\n(c) Potassium superoxide dissolves in water?

Here's what happens in each scenario:\n\n**(a) Sodium metal is dropped in water?**\nWhen sodium metal is dropped in water, it reacts vigorously and exothermically to produce sodium hydroxide and hydrogen gas. The reaction is highly exothermic, releasing enough heat to melt the sodium and often ignite the hydrogen gas, which burns with a characteristic yellow flame (due to sodium vapor). The sodium metal skims across the surface of the water due to the propulsion of the hydrogen gas.\n\n**Reaction:**\n2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) + Heat\n\n**(b) Sodium peroxide dissolves in water?**\nWhen sodium peroxide (Na2O2) dissolves in water, it undergoes hydrolysis to produce sodium hydroxide and hydrogen peroxide (H2O2). The solution becomes alkaline due to the formation of NaOH.\n\n**Reaction:**\nNa2O2(s) + 2H2O(l) → 2NaOH(aq) + H2O2(aq)\n\n**(c) Potassium superoxide dissolves in water?**\nWhen potassium superoxide (KO2) dissolves in water, it reacts to produce potassium hydroxide, hydrogen peroxide, and oxygen gas. This reaction is used in breathing apparatus as KO2 can absorb CO2 and release O2.\n\n**Reaction:**\n2KO2(s) + 2H2O(l) → 2KOH(aq) + H2O2(aq) + O2(g)

Question 1

Q1. What are the common physical and chemical features of alkali metals?

Accepted Answer

Alkali metals (Group 1 elements: Lithium, Sodium, Potassium, Rubidium, Caesium, and Francium) exhibit characteristic physical and chemical properties due to their electronic configuration (ns¹). They are highly reactive and electropositive elements.

### Common Physical Features:
1.  **Electronic Configuration:** They all have one valence electron (ns¹). This configuration makes them readily lose this electron to form a unipositive ion (M⁺).
2.  **Atomic and Ionic Radii:** They have the largest atomic and ionic radii in their respective periods. Both atomic and ionic radii increase down the group from Li to Cs due to the addition of new electron shells.
3.  **Ionization Enthalpy:** They have very low ionization enthalpies, which decrease down the group. This is because the outermost electron is far from the nucleus and is effectively shielded, making it easy to remove. The first ionization enthalpy is significantly lower than the second ionization enthalpy.
4.  **Metallic Character:** They are typical metals with a silvery-white appearance (except caesium, which is golden yellow). They are soft and can be cut with a knife, with softness increasing down the group.
5.  **Density:** They have low densities, which generally increase down the group from Li to Cs. However, potassium is lighter than sodium.
6.  **Melting and Boiling Points:** They have low melting and boiling points, which decrease down the group. This is due to weak metallic bonding resulting from the presence of only one valence electron per atom.
7.  **Metallic Lustre and Conductivity:** They possess a characteristic metallic lustre when freshly cut due to the oscillation of valence electrons. They are excellent conductors of heat and electricity.
8.  **Flame Colouration:** When heated in a Bunsen flame, their valence electrons get excited to higher energy levels. Upon returning to the ground state, they emit light of characteristic wavelengths, imparting distinct colours to the flame:
    *   Lithium (Li): Crimson red
    *   Sodium (Na): Golden yellow
    *   Potassium (K): Lilac (pale violet)
    *   Rubidium (Rb): Red-violet
    *   Caesium (Cs): Blue

### Common Chemical Features:
1.  **Reactivity:** Alkali metals are highly reactive due to their low ionization enthalpies and strong electropositive nature. Their reactivity increases down the group.
2.  **Oxidation State:** They invariably exhibit a +1 oxidation state in their compounds, as they readily lose their single valence electron to form stable M⁺ ions.
3.  **Reducing Nature:** They are strong reducing agents, meaning they readily donate electrons. Their reducing power generally increases down the group, with lithium being an exception (it is the strongest reducing agent in aqueous solution due to high hydration enthalpy).
4.  **Reaction with Air/Oxygen:** They tarnish rapidly in dry air due to the formation of an oxide layer. They react vigorously with oxygen to form different types of oxides:
    *   Lithium forms mainly monoxide (Li₂O).
    *   Sodium forms peroxide (Na₂O₂).
    *   Potassium, Rubidium, and Caesium form superoxides (MO₂).
5.  **Reaction with Water:** They react vigorously with water to form hydroxides and hydrogen gas. The reaction becomes increasingly violent down the group.
    2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)
6.  **Reaction with Halogens:** They react vigorously with halogens to form ionic halides (MX). The reactivity with halogens decreases from F₂ to I₂.
    2M(s) + X₂(g) → 2MX(s)
7.  **Reaction with Hydrogen:** They react with hydrogen at about 673 K to form ionic hydrides (MH).
    2M(s) + H₂(g) → 2MH(s)
8.  **Reaction with Liquid Ammonia:** They dissolve in liquid ammonia to give deep blue solutions, which are conducting and contain ammoniated electrons and ammoniated metal ions. These solutions are strongly reducing.
    M(s) + (x+y)NH₃(l) → [M(NH₃)ₓ]⁺ + [e(NH₃)ᵧ]⁻

Question 2

Q2. Discuss the general characteristics and gradation in properties of alkaline earth metals.

Accepted Answer

Alkaline earth metals (Group 2 elements: Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium) are less electropositive than alkali metals but are still quite reactive. Their properties are influenced by their electronic configuration (ns²). ### General Characteristics: 1. **Electronic Configuration:** All alkaline earth metals have two valence electrons in their outermost s-orbital (ns²). They tend to lose these two electrons to form dipositive ions (M²⁺). 2. **Atomic and Ionic Radii:** They have smaller atomic and ionic radii compared to their corresponding alkali metals in the same period due to a higher nuclear charge. Both atomic and ionic radii increase down the group from Be to Ba. 3. **Ionization Enthalpy:** They have higher first ionization enthalpies than alkali metals in the same period because of their smaller size and greater nuclear charge. However, their first ionization enthalpy is still relatively low, and the second ionization enthalpy is higher than the first but lower than the third. Ionization enthalpies decrease down the group. 4. **Hydration Enthalpy:** The hydration enthalpies of alkaline earth metal ions (M²⁺) are generally higher than those of alkali metal ions (M⁺) due to their smaller size and higher charge. Hydration enthalpy decreases down the group (Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺). 5. **Metallic Character:** They are silvery-white, lustrous, and relatively soft metals, though harder than alkali metals. Metallic character increases down the group. 6. **Density:** They have higher densities than alkali metals. Density generally increases down the group, but Ca is an exception (density of Ca is less than Mg). 7. **Melting and Boiling Points:** They have relatively high melting and boiling points compared to alkali metals, which generally decrease down the group (except for Mg). 8. **Oxidation State:** They invariably exhibit a +2 oxidation state in their compounds. 9. **Reducing Nature:** They are strong reducing agents, though less powerful than alkali metals. Their reducing power generally increases down the group. 10. **Flame Colouration:** Except for Beryllium and Magnesium, the alkaline earth metals impart characteristic colours to the Bunsen flame due to the excitation and de-excitation of their valence electrons: * Calcium (Ca): Brick red * Strontium (Sr): Crimson red * Barium (Ba): Apple green ### Gradation in Properties (Trends down the group from Be to Ba): 1. **Electronic Configuration:** Remains ns² for all elements. 2. **Atomic and Ionic Radii:** Increase down the group due to the addition of new electron shells, leading to increased shielding and greater distance of valence electrons from the nucleus. 3. **Ionization Enthalpy:** Decreases down the group. As the atomic size increases, the outermost electrons are further from the nucleus and experience less attraction, making them easier to remove. 4. **Hydration Enthalpy:** Decreases down the group. As the ionic size increases, the charge density decreases, leading to weaker interaction with water molecules. 5. **Metallic Character:** Increases down the group. The ease of losing electrons increases with increasing atomic size and decreasing ionization enthalpy. 6. **Density:** Generally increases down the group, but with some irregularities (e.g., Ca is less dense than Mg). This trend is influenced by both atomic mass and crystal structure. 7. **Melting and Boiling Points:** Generally decrease down the group, though not as smoothly as alkali metals (e.g., Mg has a lower melting point than Ca). This is due to variations in crystal structure and metallic bonding strength. 8. **Reactivity:** Increases down the group. The ease of losing electrons (lower ionization enthalpy) makes them more reactive. 9. **Basic Character of Oxides and Hydroxides:** Increases down the group. As the metallic character increases, the basicity of their oxides (MO) and hydroxides (M(OH)₂) increases. BeO and Be(OH)₂ are amphoteric, while others are basic. 10. **Solubility of Sulphates:** Decreases down the group (BeSO₄ > MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄). This is due to the decreasing hydration enthalpy of the M²⁺ ions outweighing the decreasing lattice enthalpy. 11. **Thermal Stability of Carbonates:** Increases down the group (BeCO₃ < MgCO₃ < CaCO₃ < SrCO₃ < BaCO₃). Larger cations stabilize larger anions more effectively, and the polarizing power of the cation decreases down the group.

Question 3

Q3. Why are alkali metals not found in free state in nature?

Accepted Answer

Alkali metals are not found in the free state in nature because they are highly reactive and electropositive elements. This high reactivity stems from two main factors:

1.  **Low Ionization Enthalpy:** Alkali metals have only one electron in their outermost shell (ns¹). This electron is loosely held by the nucleus and can be easily removed. Consequently, they have very low ionization enthalpies, meaning little energy is required to form a positive ion (M⁺).
2.  **Strong Electropositive Character:** Due to their low ionization enthalpy, alkali metals have a strong tendency to lose their valence electron and form stable positive ions. This makes them highly electropositive.

Because of their strong tendency to lose electrons, alkali metals readily react with other elements (like oxygen, halogens, sulfur, etc.) present in their surroundings to form stable ionic compounds. They are never found as free metals but always exist in combined forms, such as chlorides (e.g., NaCl, KCl), oxides, silicates, and carbonates.

Question 4

Q4. How would you explain the reducing power of alkali metals?

Accepted Answer

Alkali metals are strong reducing agents. A reducing agent is a substance that readily loses electrons (gets oxidized) and causes another substance to gain electrons (get reduced).

The strong reducing power of alkali metals can be explained by the following factors:

1.  **Low Ionization Enthalpy:** Alkali metals have very low ionization enthalpies. They have only one electron in their outermost shell (ns¹), which is loosely held and can be easily removed to form a stable M⁺ ion. The ease with which they lose an electron makes them excellent electron donors, hence strong reducing agents.

2.  **Highly Negative Standard Electrode Potentials (E°):** The reducing power of a metal in aqueous solution is quantitatively measured by its standard electrode potential (E°). A more negative E° value indicates a stronger reducing agent. For alkali metals, the overall process for reduction is:
    M(s) → M⁺(aq) + e⁻
    This process involves three steps:
    a.  **Sublimation Enthalpy (ΔH_sub):** M(s) → M(g) (Energy required to convert solid metal to gaseous atom)
    b.  **Ionization Enthalpy (ΔH_ion):** M(g) → M⁺(g) + e⁻ (Energy required to remove an electron from a gaseous atom)
    c.  **Hydration Enthalpy (ΔH_hyd):** M⁺(g) + H₂O → M⁺(aq) (Energy released when the gaseous ion is hydrated)

The sum of these enthalpy changes determines the overall energy change for the formation of the hydrated ion. A more negative overall energy change (more exothermic) favors the formation of M⁺(aq) and thus indicates a stronger reducing power. Alkali metals have low ionization enthalpies and significant hydration enthalpies (especially for smaller ions like Li⁺), contributing to their highly negative standard electrode potentials.

**Trend in Reducing Power:** Generally, reducing power increases down the group as ionization enthalpy decreases. However, **Lithium (Li)** is an exception. Despite having a higher ionization enthalpy than other alkali metals, it is the strongest reducing agent in aqueous solution. This is because of its exceptionally high hydration enthalpy due to the very small size of the Li⁺ ion. The large amount of energy released during the hydration of Li⁺ compensates for its higher ionization enthalpy, making its overall electrode potential the most negative among alkali metals.

Question 5

Q5. Arrange the following metals in the increasing order of their reducing power: Li, Na, K, Rb, Cs.

Accepted Answer

The increasing order of reducing power for the given alkali metals is:

**Na < K < Rb < Cs < Li**

**Explanation:**

1.  **General Trend:** For alkali metals, the reducing power generally increases down the group from Na to Cs. This is because as we move down the group, the atomic size increases, and the ionization enthalpy decreases. A lower ionization enthalpy means the electron can be removed more easily, making the element a stronger reducing agent.
    So, the expected order based on ionization enthalpy would be Na < K < Rb < Cs.

2.  **Anomalous Behavior of Lithium:** Lithium is an exception to this general trend. Despite having the highest ionization enthalpy among the alkali metals, it is the strongest reducing agent in aqueous solution. This is attributed to the exceptionally high hydration enthalpy of the small Li⁺ ion. The small size of Li⁺ leads to a very high charge density, causing it to be extensively hydrated in aqueous solution. The large amount of energy released during the hydration of Li⁺ (ΔH_hyd) more than compensates for its higher ionization enthalpy (ΔH_ion) and sublimation enthalpy (ΔH_sub), resulting in the most negative standard electrode potential (E°) for Lithium.

The overall process for reducing power in aqueous solution involves:
    M(s) → M(g) (Sublimation)
    M(g) → M⁺(g) + e⁻ (Ionization)
    M⁺(g) + H₂O → M⁺(aq) (Hydration)

The sum of these energy changes determines the standard electrode potential (E°), which is a measure of reducing power. For Lithium, the very large negative hydration enthalpy makes its E° value the most negative, indicating the strongest reducing power.

Therefore, considering both the general trend and the anomaly of Lithium, the correct increasing order of reducing power is Na < K < Rb < Cs < Li.

Question 6

Q6. Why are alkali metals not found in free state in nature?

Accepted Answer

Alkali metals are not found in the free state in nature because they are highly reactive. They have a strong tendency to lose their single valence electron (ns¹) to form a stable cation (M⁺). This low ionization enthalpy and high electropositivity make them readily react with various elements like oxygen, halogens, and water, forming stable compounds. Due to their high reactivity, they always exist in combined forms (compounds) in nature, such as halides, oxides, silicates, etc.

Question 7

Q7. Describe the general characteristics of alkali metals with respect to their electronic configuration, oxidation states, atomic and ionic radii, ionisation enthalpy and other proper

Accepted Answer

Alkali metals (Group 1 elements: Li, Na, K, Rb, Cs, Fr) are highly reactive metals with distinct general characteristics due to their electronic configuration.

### 1. Electronic Configuration
All alkali metals have one electron in their outermost s-orbital (ns¹). This general electronic configuration is responsible for their characteristic properties. For example:
*   Lithium (Li): [He] 2s¹
*   Sodium (Na): [Ne] 3s¹
*   Potassium (K): [Ar] 4s¹
*   Rubidium (Rb): [Kr] 5s¹
*   Caesium (Cs): [Xe] 6s¹
*   Francium (Fr): [Rn] 7s¹

### 2. Oxidation States
Due to the presence of a single valence electron and very low ionization enthalpies, alkali metals readily lose this electron to form a unipositive ion (M⁺). Thus, their characteristic and only stable oxidation state in compounds is +1. They do not exhibit higher oxidation states because the removal of a second electron would require a very high amount of energy (second ionization enthalpy is very high).

### 3. Atomic and Ionic Radii
Alkali metals have the largest atomic radii in their respective periods. This is because they have only one electron in their outermost shell, and the effective nuclear charge experienced by this electron is relatively low. As we move down the group from Li to Cs, the atomic and ionic radii increase due to the addition of new electron shells, which increases the distance of the valence electron from the nucleus and enhances shielding effect.

### 4. Ionisation Enthalpy
Alkali metals have very low first ionization enthalpies. This is due to their large atomic size and the presence of a single electron in the outermost shell, which is loosely held by the nucleus. As we move down the group, the first ionization enthalpy decreases from Li to Cs. This is because the atomic size increases, and the outermost electron is further away from the nucleus, making it easier to remove. The second ionization enthalpy of alkali metals is extremely high because after losing one electron, they attain a stable noble gas configuration, and removing an electron from this stable configuration requires a large amount of energy.

### 5. Other Properties
*   **Metallic Character:** They are typical metals—soft, silvery-white, and highly lustrous when freshly cut. They are good conductors of heat and electricity. Metallic character increases down the group.
*   **Density:** They have low densities, which generally increase down the group from Li to Cs (Potassium is an exception, being lighter than sodium). This is due to their large atomic volumes.
*   **Melting and Boiling Points:** They have low melting and boiling points, which decrease down the group. This is attributed to the weak metallic bonding due to the presence of only one valence electron per atom.
*   **Electropositivity/Metallic Character:** They are highly electropositive, meaning they have a strong tendency to lose electrons and form positive ions. Electropositivity increases down the group.
*   **Reducing Nature:** They are strong reducing agents due to their low ionization enthalpies and high tendency to lose electrons. Their reducing power generally increases down the group (Lithium is an exception in aqueous solution, being the strongest reducing agent).
*   **Flame Colouration:** When heated in a flame, their valence electrons get excited to higher energy levels. When these excited electrons return to their ground state, they emit light of characteristic wavelengths, resulting in distinct colours:
    *   Lithium (Li): Crimson red
    *   Sodium (Na): Golden yellow
    *   Potassium (K): Lilac (pale violet)
    *   Rubidium (Rb): Red-violet
    *   Caesium (Cs): Blue
*   **Reactivity:** They are highly reactive due to their low ionization enthalpy and high electropositivity. Reactivity increases down the group. They react vigorously with air, water, and halogens.

Question 8

Q8. Discuss the general characteristics and gradation in properties of alkaline earth metals with respect to their electronic configuration, oxidation states, atomic and ionic radii, i

Accepted Answer

Alkaline earth metals (Group 2 elements: Be, Mg, Ca, Sr, Ba, Ra) are less reactive than alkali metals but still quite reactive. Their properties are largely determined by their electronic configuration.

### 1. Electronic Configuration
All alkaline earth metals have two electrons in their outermost s-orbital (ns²). This general electronic configuration is responsible for their characteristic properties.
*   Beryllium (Be): [He] 2s²
*   Magnesium (Mg): [Ne] 3s²
*   Calcium (Ca): [Ar] 4s²
*   Strontium (Sr): [Kr] 5s²
*   Barium (Ba): [Xe] 6s²
*   Radium (Ra): [Rn] 7s²

### 2. Oxidation States
Due to the presence of two valence electrons and relatively low first and second ionization enthalpies, alkaline earth metals readily lose both electrons to form a dipositive ion (M²⁺). Thus, their characteristic and only stable oxidation state in compounds is +2. The sum of the first and second ionization enthalpies is relatively low, making the formation of M²⁺ ions energetically favorable. The third ionization enthalpy is very high, preventing the formation of M³⁺ ions.

### 3. Atomic and Ionic Radii
Alkaline earth metals have smaller atomic radii compared to their corresponding alkali metals in the same period. This is because they have a higher nuclear charge (+2 vs +1) which pulls the electrons closer to the nucleus. As we move down the group from Be to Ba, the atomic and ionic radii increase due to the addition of new electron shells, which increases the distance of the valence electrons from the nucleus and enhances the shielding effect.

### 4. Ionisation Enthalpy
Alkaline earth metals have higher first ionization enthalpies compared to alkali metals in the same period due to their smaller atomic size and higher effective nuclear charge. However, their first ionization enthalpies are still relatively low. The second ionization enthalpy is also low enough to allow the formation of M²⁺ ions. As we move down the group, both the first and second ionization enthalpies decrease from Be to Ba. This is because the atomic size increases, and the outermost electrons are further away from the nucleus, making them easier to remove. The third ionization enthalpy is very high, as removing an electron from the stable M²⁺ ion (which has a noble gas configuration) requires a large amount of energy.

### 5. Other Properties
*   **Metallic Character:** They are silvery-white, lustrous, and relatively soft metals, though harder than alkali metals. They are good conductors of heat and electricity. Metallic character increases down the group.
*   **Density:** They have higher densities than alkali metals in the same period. Density generally increases down the group from Be to Ba (Magnesium is an exception, being lighter than calcium). This is due to their smaller atomic volumes and higher atomic masses.
*   **Melting and Boiling Points:** They have higher melting and boiling points compared to alkali metals. This is attributed to stronger metallic bonding due to the presence of two valence electrons per atom. Melting and boiling points do not show a regular trend, with Be and Mg having relatively high values, and Ca, Sr, Ba showing a decrease.
*   **Electropositivity/Metallic Character:** They are highly electropositive, though less so than alkali metals. Electropositivity increases down the group.
*   **Reducing Nature:** They are strong reducing agents, though less powerful than alkali metals. Their reducing power generally increases down the group.
*   **Flame Colouration:** When heated in a flame, their valence electrons get excited to higher energy levels. When these excited electrons return to their ground state, they emit light of characteristic wavelengths, resulting in distinct colours. Beryllium and Magnesium do not impart colour to the flame because their electrons are too strongly bound to be excited by the flame energy.
    *   Calcium (Ca): Brick red
    *   Strontium (Sr): Crimson red
    *   Barium (Ba): Apple green
*   **Reactivity:** They are reactive, though less reactive than alkali metals. Reactivity generally increases down the group. They react with air, water, and halogens, but less vigorously than alkali metals.

Question 9

Q9. Why are alkali metals strong reducing agents?

Accepted Answer

Alkali metals are strong reducing agents because they have a very high tendency to lose their single valence electron (ns¹) to form a stable unipositive ion (M⁺) with a noble gas configuration. This tendency is due to:

1.  **Low Ionization Enthalpy:** They have the lowest first ionization enthalpies in their respective periods. This means only a small amount of energy is required to remove their outermost electron.
2.  **Large Atomic Size:** Their large atomic size results in the valence electron being far from the nucleus and experiencing less effective nuclear charge, making it easier to remove.
3.  **High Electropositivity:** Their strong tendency to lose electrons makes them highly electropositive.

By losing an electron, alkali metals themselves get oxidized (M → M⁺ + e⁻) and in turn, reduce other species. The ease with which they get oxidized makes them powerful reducing agents. Lithium, despite having a higher ionization enthalpy than other alkali metals, is the strongest reducing agent in aqueous solution due to its exceptionally high hydration enthalpy of the small Li⁺ ion.

Question 10

Q10. When an alkali metal dissolves in liquid ammonia the solution acquires different colours. Explain this type of behaviour.

Accepted Answer

When an alkali metal (M) dissolves in liquid ammonia, it forms a deep blue solution. This unique behaviour is due to the formation of ammoniated electrons and ammoniated metal ions.

**1. Dissolution Process:**
Upon dissolution, the alkali metal atom readily loses its valence electron to form a metal cation (M⁺) and an electron (e⁻). Both the cation and the electron get solvated (ammoniated) by ammonia molecules.

**Equation:**
M(s) + (x + y) NH₃(l) → [M(NH₃)ₓ]⁺ (ammoniated cation) + [e(NH₃)y]⁻ (ammoniated electron)

**2. Deep Blue Colour:**
The deep blue colour of the solution is primarily due to the presence of the **ammoniated electrons** ([e(NH₃)y]⁻). These electrons absorb energy in the visible region of the electromagnetic spectrum, getting excited to higher energy levels. When they return to their ground state, they emit light corresponding to the blue colour, or rather, absorb light in the red/orange region, transmitting blue light. The solutions are also paramagnetic due to the presence of unpaired ammoniated electrons.

**3. Conductivity:**
The solutions are highly electrically conductive, similar to metals, due to the presence of both ammoniated cations and ammoniated electrons.

**4. Colour Change at Higher Concentrations:**
At higher concentrations (above 3 M), the blue solutions gradually change to a bronze or copper-like colour and become diamagnetic. This is because the ammoniated electrons start to associate or pair up, forming electron clusters. The metallic lustre is due to the formation of these electron clusters, which reflect light like a metal. The diamagnetic nature arises from the pairing of electrons.

**5. Stability and Decomposition:**
The solutions are metastable and slowly decompose to form metal amides and hydrogen gas, especially in the presence of impurities or catalysts:
2M(s) + 2NH₃(l) → 2MNH₂(ammoniated amide) + H₂(g)
This decomposition is faster for lithium.

Question 11

Q11. Why are alkali metals not found in free state in nature?

Accepted Answer

Alkali metals are highly reactive elements due to their electronic configuration and low ionisation enthalpies. They have only one electron in their outermost shell, which they readily lose to form a stable unipositive ion (M⁺). This tendency to lose an electron makes them strong reducing agents and highly electropositive. Consequently, they react vigorously with various elements like oxygen, halogens, and water present in the environment to form stable ionic compounds. Because of their high reactivity and strong tendency to form compounds, they are never found in the free or elemental state in nature; instead, they exist only in combined forms as salts, oxides, silicates, etc.

Question 12

Q12. Describe the general characteristics of alkali metals with respect to their electronic configuration, oxidation states, atomic and ionic radii, ionisation enthalpy and other proper

Accepted Answer

Alkali metals are the elements of Group 1 of the periodic table, comprising lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). They are characterised by the following general properties:

### 1. Electronic Configuration
All alkali metals have one valence electron (ns¹) in their outermost shell. This general electronic configuration is responsible for their characteristic properties.
*   Li: [He] 2s¹
*   Na: [Ne] 3s¹
*   K: [Ar] 4s¹
*   Rb: [Kr] 5s¹
*   Cs: [Xe] 6s¹
*   Fr: [Rn] 7s¹

### 2. Oxidation States
Due to the presence of a single valence electron (ns¹), alkali metals readily lose this electron to achieve a stable noble gas configuration. Therefore, they invariably exhibit a +1 oxidation state in their compounds. They do not show higher oxidation states because the removal of a second electron would require a very high amount of energy (due to stable noble gas core).

### 3. Atomic and Ionic Radii
Alkali metals have the largest atomic radii in their respective periods because they have the lowest effective nuclear charge. Down the group, both atomic and ionic radii increase progressively due to the addition of new electron shells. The ionic radii (M⁺) are significantly smaller than their corresponding atomic radii (M) because of the loss of the outermost electron and an increase in the effective nuclear charge on the remaining electrons.

### 4. Ionisation Enthalpy
Alkali metals have very low ionisation enthalpies (IE₁) because of their large atomic size and the ease with which they can lose their single valence electron. As we move down the group, the ionisation enthalpy decreases from Li to Cs. This is because the atomic size increases, and the outermost electron is further away from the nucleus, experiencing less attraction, thus requiring less energy to remove. The second ionisation enthalpy (IE₂) is extremely high for alkali metals, as it involves removing an electron from a stable noble gas core configuration.

### 5. Other Properties
*   **Metallic Character:** They are typical metals, exhibiting metallic lustre, high electrical and thermal conductivity. These properties are due to the presence of delocalised electrons.
*   **Density:** They have low densities, which generally increase down the group from Li to Cs. However, potassium is lighter than sodium due to its unusually large atomic volume.
*   **Melting and Boiling Points:** Alkali metals have low melting and boiling points. This is attributed to their weak metallic bonding, which arises from the presence of only one valence electron per atom and their large atomic size. The melting and boiling points generally decrease down the group.
*   **Electropositivity:** They are highly electropositive elements, meaning they have a strong tendency to lose electrons and form positive ions. This electropositivity increases down the group.
*   **Reducing Nature:** Due to their low ionisation enthalpies and high electropositivity, alkali metals are powerful reducing agents. Their reducing power generally increases down the group, with lithium being an exception (it is the strongest reducing agent in aqueous solution due to high hydration enthalpy).
*   **Flame Colouration:** When heated in a Bunsen flame, alkali metals and their salts impart characteristic colours to the flame. This is because the heat from the flame excites the outermost electron to a higher energy level. When the excited electron returns to its ground state, it emits radiation in the visible region of the spectrum.
    *   Li: Crimson red
    *   Na: Golden yellow
    *   K: Lilac (pale violet)
    *   Rb: Red-violet
    *   Cs: Blue
*   **Reactivity:** They are highly reactive elements, reacting vigorously with air (oxygen), water, and halogens. Their reactivity increases down the group due to decreasing ionisation enthalpy.
*   **Hydration Enthalpy:** The hydration enthalpy of alkali metal ions (M⁺) decreases with increasing ionic size. Li⁺ has the highest hydration enthalpy, leading to its high reducing power in aqueous solutions.

Question 13

Q13. Discuss the anomalous behaviour of lithium with respect to other members of the Group 1 elements.

Accepted Answer

Lithium, the first element of Group 1, exhibits anomalous behaviour compared to the other members of the alkali metal family (Na, K, Rb, Cs). This anomalous behaviour is primarily due to:
1.  **Smallest atomic and ionic size:** Lithium has the smallest atomic and ionic radii among alkali metals.
2.  **High ionisation enthalpy:** Due to its small size, the valence electron is held more tightly, resulting in a relatively higher ionisation enthalpy.
3.  **High electronegativity:** It has the highest electronegativity among alkali metals.
4.  **High polarising power:** The Li⁺ ion has a very high charge density (charge/radius ratio), which gives it a high tendency to polarise electron clouds of anions, leading to a greater covalent character in its compounds.
5.  **Absence of d-orbitals:** Unlike heavier elements, lithium does not have vacant d-orbitals in its valence shell, limiting its maximum covalency to four.

These factors lead to the following distinct properties of lithium:

### 1. Hardness and Melting Point
Lithium is much harder than other alkali metals and has a higher melting and boiling point. This is due to stronger metallic bonding resulting from its small size and higher cohesive energy.

### 2. Reactivity with Air/Oxygen
*   Lithium reacts with oxygen to primarily form lithium oxide (Li₂O), and only a small amount of peroxide. Other alkali metals form peroxides (Na₂O₂) and superoxides (KO₂, RbO₂, CsO₂).
    *   4Li(s) + O₂(g) → 2Li₂O(s)
*   Lithium is the only alkali metal that reacts directly with nitrogen to form lithium nitride (Li₃N) due to the high lattice energy of Li₃N. Other alkali metals do not react with nitrogen under normal conditions.
    *   6Li(s) + N₂(g) → 2Li₃N(s)

### 3. Reactivity with Water
Lithium reacts less vigorously with water than other alkali metals. While still reactive, its reaction is less explosive compared to sodium or potassium. This is due to its higher hydration enthalpy and higher electrode potential.

### 4. Formation of Hydrides
Lithium forms a more stable hydride (LiH) than other alkali metals.

### 5. Solubility of Compounds
*   Lithium salts of large anions (e.g., LiClO₄, LiF, Li₂CO₃) are less soluble in water than the corresponding salts of other alkali metals. This is due to the high lattice energy of lithium compounds.
*   Conversely, lithium salts of small anions (e.g., LiCl, LiNO₃) are more soluble in organic solvents (like ethanol) due to their greater covalent character.

### 6. Thermal Stability of Carbonates and Nitrates
*   Lithium carbonate (Li₂CO₃) is thermally less stable and decomposes easily on heating to form lithium oxide and carbon dioxide. Other alkali metal carbonates are stable to heat.
    *   Li₂CO₃(s) → Li₂O(s) + CO₂(g)
*   Lithium nitrate (LiNO₃) decomposes on heating to give lithium oxide, nitrogen dioxide, and oxygen. Other alkali metal nitrates decompose to give the nitrite and oxygen.
    *   4LiNO₃(s) → 2Li₂O(s) + 4NO₂(g) + O₂(g)
    *   2NaNO₃(s) → 2NaNO₂(s) + O₂(g)

### 7. Formation of Complex Compounds
Due to its small size and high charge density, Li⁺ has a greater tendency to form complex compounds than other alkali metal ions.

### 8. Nature of Hydroxide
Lithium hydroxide (LiOH) is a weaker base than other alkali metal hydroxides and is less soluble in water. It is also less stable thermally.

### 9. Diagonal Relationship with Magnesium
Lithium shows a diagonal relationship with magnesium (Group 2, Period 3), exhibiting several similarities in properties, which is another aspect of its anomalous behaviour within its own group.

Question 14

Q14. What are the similarities between lithium and magnesium?

Accepted Answer

Lithium (Group 1, Period 2) and magnesium (Group 2, Period 3) exhibit a diagonal relationship, meaning they show similar properties despite being in different groups. This similarity arises because their ionic sizes and charge/radius ratios are very similar (Li⁺: 0.76 Å, Mg²⁺: 0.72 Å). The similarities are as follows:

### 1. Atomic and Ionic Sizes
Both Li⁺ and Mg²⁺ ions have very similar ionic radii (Li⁺ ≈ 76 pm, Mg²⁺ ≈ 72 pm), leading to comparable charge densities and polarising powers.

### 2. Hardness
Both lithium and magnesium are relatively hard metals compared to other elements in their respective groups.

### 3. Melting and Boiling Points
Both have relatively high melting and boiling points compared to other alkali metals (for Li) and alkaline earth metals (for Mg, though less pronounced).

### 4. Reactivity with Oxygen
*   Both react with oxygen to form primarily normal oxides (Li₂O and MgO) and do not form peroxides or superoxides readily under normal conditions (unlike Na which forms Na₂O₂ and K which forms KO₂).
    *   4Li(s) + O₂(g) → 2Li₂O(s)
    *   2Mg(s) + O₂(g) → 2MgO(s)

### 5. Reactivity with Nitrogen
Both lithium and magnesium are the only elements in their respective groups that react directly with nitrogen at room temperature (or on heating) to form nitrides.
*   6Li(s) + N₂(g) → 2Li₃N(s)
*   3Mg(s) + N₂(g) → Mg₃N₂(s)

### 6. Formation of Carbides
Both lithium and magnesium form carbides (Li₂C₂ and MgC₂) which react with water to produce acetylene (C₂H₂).
*   Li₂C₂(s) + 2H₂O(l) → 2LiOH(aq) + C₂H₂(g)
*   Mg₂C₂(s) + 2H₂O(l) → 2Mg(OH)₂(aq) + C₂H₂(g)

### 7. Hydroxides
Both lithium hydroxide (LiOH) and magnesium hydroxide (Mg(OH)₂) are relatively weak bases and are sparingly soluble in water. They decompose on heating.
*   2LiOH(s) → Li₂O(s) + H₂O(g)
*   Mg(OH)₂(s) → MgO(s) + H₂O(g)

### 8. Carbonates
Both lithium carbonate (Li₂CO₃) and magnesium carbonate (MgCO₃) are thermally unstable and decompose on heating to form their respective oxides and carbon dioxide.
*   Li₂CO₃(s) → Li₂O(s) + CO₂(g)
*   MgCO₃(s) → MgO(s) + CO₂(g)

### 9. Bicarbonates
Both lithium and magnesium do not form solid bicarbonates. They exist only in solution.

### 10. Chlorides
Both lithium chloride (LiCl) and magnesium chloride (MgCl₂) are deliquescent (absorb moisture from the air) and crystallise from aqueous solution as hydrates (LiCl·2H₂O and MgCl₂·6H₂O). They are also soluble in organic solvents like ethanol and pyridine due to their significant covalent character.

### 11. Formation of Complex Compounds
Both Li⁺ and Mg²⁺ ions have a greater tendency to form complex compounds compared to other members of their respective groups, owing to their small size and high charge density.

### 12. Electronegativity
Their electronegativity values are quite similar (Li: 0.98, Mg: 1.31 on Pauling scale), which contributes to the covalent character of their compounds.

Question 15

Q15. Explain why alkali metals have low melting and boiling points.

Accepted Answer

Alkali metals have low melting and boiling points primarily due to the nature of their metallic bonding. Here's a detailed explanation:

1.  **Weak Metallic Bonding:** Alkali metals have only one valence electron per atom (ns¹). These single valence electrons are delocalised throughout the metallic lattice, forming a 'sea of electrons' that binds the positively charged metal ions (M⁺) together. However, because there is only one electron per atom contributing to this metallic bond, the overall metallic bonding forces are relatively weak.

2.  **Large Atomic Size:** Alkali metals have the largest atomic radii in their respective periods. The larger the atomic size, the greater the distance between the nucleus and the delocalised valence electrons. This increased distance results in weaker electrostatic attraction between the positive metal ions and the electron sea.

3.  **Low Cohesive Energy:** The combination of weak metallic bonding and large atomic size leads to low cohesive energy within the metallic lattice. Cohesive energy is the energy required to separate the atoms in a solid.

To melt or boil a metal, sufficient energy must be supplied to overcome these metallic bonds. Since the metallic bonds in alkali metals are weak, only a small amount of thermal energy is required to break them, resulting in low melting and boiling points. The melting and boiling points generally decrease down the group (from Li to Cs) as the atomic size increases further, leading to even weaker metallic bonding.

Question 16

Q16. Why are lithium salts commonly hydrated and those of other alkali ions usually anhydrous?

Accepted Answer

Lithium salts are commonly hydrated due to the exceptionally small size and high charge density of the Li+ ion. This leads to a very high charge-to-radius ratio and consequently, a very high hydration enthalpy. The strong electrostatic attraction between the small Li+ ion and the polar water molecules results in the formation of stable hydrated salts, such as LiCl·2H2O.

In contrast, other alkali metal ions (Na+, K+, Rb+, Cs+) are much larger in size and have significantly lower charge densities. Their hydration enthalpies are much smaller and are often not sufficient to overcome the lattice energy to form stable hydrates. Therefore, their salts are usually anhydrous.

Question 17

Q17. Why is LiF almost insoluble in water whereas LiCl is soluble not only in water but also in acetone?

Accepted Answer

The solubility of an ionic compound in water depends on the balance between its lattice energy and hydration energy. For a compound to be soluble, the hydration energy must be greater than or comparable to the lattice energy.

**LiF is almost insoluble in water:**
LiF has a very high lattice energy. Both Li+ and F- ions are exceptionally small. The strong electrostatic forces of attraction between these small, oppositely charged ions result in a very compact crystal lattice with high lattice energy. Although Li+ has a high hydration energy, the lattice energy of LiF is so high that it cannot be overcome by the hydration energy, making LiF sparingly soluble or almost insoluble in water.

**LiCl is soluble in water and acetone:**
1.  **Solubility in water:** LiCl has a lower lattice energy compared to LiF because the Cl- ion is much larger than the F- ion. While the hydration energy of Li+ is still high, it is sufficient to overcome the relatively lower lattice energy of LiCl, making it soluble in water.
2.  **Solubility in acetone:** LiCl exhibits significant covalent character. According to Fajan's rules, the small size and high charge density of Li+ ion can polarize the larger electron cloud of the Cl- ion to a greater extent, leading to increased covalent character in the Li-Cl bond. Acetone is a polar organic solvent. The covalent character of LiCl allows it to dissolve in organic solvents like acetone (which is also polar) based on the principle of 'like dissolves like'.

Question 18

Q18. Explain the significance of the terms 'biologically important' in context of Na+, K+, Mg2+ and Ca2+.

Accepted Answer

The terms 'biologically important' for Na+, K+, Mg2+, and Ca2+ refer to their essential roles in maintaining various physiological functions, biochemical processes, and structural integrity within living organisms. These ions are crucial for life and their imbalance can lead to severe health issues.

### 1. Sodium Ions (Na+)
Sodium ions are the most abundant cations in the extracellular fluid (outside cells) and play a vital role in:
*   **Fluid and Electrolyte Balance:** Along with chloride ions, Na+ helps maintain osmotic pressure and water balance between cells and the extracellular fluid.
*   **Nerve Impulse Transmission:** The differential concentration of Na+ and K+ across the nerve cell membrane is critical for generating and transmitting nerve impulses (action potentials) via the Na+/K+ pump.
*   **Glucose and Amino Acid Absorption:** Na+ cotransport systems are involved in the absorption of glucose and amino acids from the intestine and their reabsorption in the kidneys.
*   **Muscle Contraction:** Na+ influx is part of the depolarization phase in muscle contraction.

### 2. Potassium Ions (K+)
Potassium ions are the most abundant cations within the intracellular fluid (inside cells) and are essential for:
*   **Nerve Impulse Transmission:** K+ plays a crucial role in repolarization of nerve cells after an action potential, working in conjunction with Na+.
*   **Muscle Contraction:** K+ is vital for the proper functioning of muscles, including the heart muscle. Imbalances can lead to cardiac arrhythmias.
*   **Enzyme Activation:** K+ acts as a cofactor for several enzymes involved in metabolism.
*   **Maintaining Cell Volume:** K+ contributes to the osmotic balance and maintenance of cell volume.

### 3. Magnesium Ions (Mg2+)
Magnesium ions are the second most abundant intracellular cations and are involved in over 300 enzymatic reactions, making them critical for:
*   **Enzyme Cofactor:** Mg2+ is an essential cofactor for enzymes involved in energy production (e.g., ATPases, enzymes in glycolysis and oxidative phosphorylation), nucleic acid synthesis (DNA and RNA polymerase), protein synthesis, and fatty acid synthesis.
*   **Bone and Teeth Formation:** A significant portion of the body's magnesium is found in bones, contributing to their structure and strength.
*   **Muscle and Nerve Function:** Mg2+ helps regulate muscle contraction and nerve impulse transmission. It acts as a natural calcium channel blocker, promoting muscle relaxation.
*   **Chlorophyll Component:** In plants, Mg2+ is the central metal ion in the chlorophyll molecule, essential for photosynthesis.

### 4. Calcium Ions (Ca2+)
Calcium ions are the most abundant mineral in the human body and are indispensable for:
*   **Bone and Teeth Formation:** Ca2+ is the primary structural component of bones and teeth, providing strength and rigidity.
*   **Muscle Contraction:** Ca2+ ions are released from the sarcoplasmic reticulum and bind to troponin, initiating the cascade of events leading to muscle contraction.
*   **Blood Clotting:** Ca2+ is an essential cofactor for several enzymes (clotting factors) in the blood coagulation cascade.
*   **Nerve Impulse Transmission:** Ca2+ influx into nerve terminals triggers the release of neurotransmitters.
*   **Hormone Secretion:** Ca2+ acts as a second messenger in many cellular signaling pathways, regulating the secretion of hormones and enzymes.
*   **Cardiac Function:** Ca2+ is crucial for the rhythmic beating of the heart.

Question 19

Q19. What happens when
(a) Sodium metal is dropped in water?
(b) Sodium peroxide dissolves in water?
(c) Potassium superoxide dissolves in water?

Accepted Answer

Here's what happens in each scenario:

**(a) Sodium metal is dropped in water?**
When sodium metal is dropped in water, it reacts vigorously and exothermically to produce sodium hydroxide and hydrogen gas. The reaction is highly exothermic, releasing enough heat to melt the sodium and often ignite the hydrogen gas, which burns with a characteristic yellow flame (due to sodium vapor). The sodium metal skims across the surface of the water due to the propulsion of the hydrogen gas.

**Reaction:**
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) + Heat

**(b) Sodium peroxide dissolves in water?**
When sodium peroxide (Na2O2) dissolves in water, it undergoes hydrolysis to produce sodium hydroxide and hydrogen peroxide (H2O2). The solution becomes alkaline due to the formation of NaOH.

**Reaction:**
Na2O2(s) + 2H2O(l) → 2NaOH(aq) + H2O2(aq)

**(c) Potassium superoxide dissolves in water?**
When potassium superoxide (KO2) dissolves in water, it reacts to produce potassium hydroxide, hydrogen peroxide, and oxygen gas. This reaction is used in breathing apparatus as KO2 can absorb CO2 and release O2.

**Reaction:**
2KO2(s) + 2H2O(l) → 2KOH(aq) + H2O2(aq) + O2(g)

Question 20

Q20. Compare the solubility and thermal stability of the following compounds of the alkali metals with those of the alkaline earth metals.
(a) Nitrates
(b) Carbonates
(c) Sulphates

Accepted Answer

The solubility and thermal stability of compounds of alkali metals (Group 1) and alkaline earth metals (Group 2) show distinct trends due to differences in ionic size, charge, lattice energy, and polarizing power. ### General Comparison: * **Alkali Metals (AM):** Larger ionic size, +1 charge, lower charge density, lower polarizing power. * **Alkaline Earth Metals (AEM):** Smaller ionic size (for comparable period), +2 charge, higher charge density, higher polarizing power. ### (a) Nitrates (MNO3 and M'(NO3)2) **1. Solubility:** * **Alkali Metal Nitrates:** All alkali metal nitrates (e.g., NaNO3, KNO3) are highly soluble in water. This is because the hydration energy of the alkali metal ions is generally sufficient to overcome their relatively lower lattice energies. * **Alkaline Earth Metal Nitrates:** Most alkaline earth metal nitrates (e.g., Mg(NO3)2, Ca(NO3)2, Sr(NO3)2, Ba(NO3)2) are also soluble in water. However, their solubility is generally lower than that of alkali metal nitrates due to higher lattice energies (resulting from the +2 charge on the cation). Within the group, solubility generally decreases down the group for larger anions, but for nitrates, it tends to remain high or slightly decrease. **2. Thermal Stability:** * **Alkali Metal Nitrates:** Alkali metal nitrates are generally more thermally stable than alkaline earth metal nitrates. They decompose at higher temperatures. Lithium nitrate (LiNO3) is an exception, decomposing like AEM nitrates. Other alkali metal nitrates decompose to nitrites and oxygen: 2MNO3(s) → 2MNO2(s) + O2(g) (M = Na, K, Rb, Cs) * **Alkaline Earth Metal Nitrates:** Alkaline earth metal nitrates are less thermally stable. Due to the higher polarizing power of the M2+ ions, they decompose at lower temperatures to form metal oxides, nitrogen dioxide, and oxygen. Lithium nitrate also decomposes similarly: 2M'(NO3)2(s) → 2M'O(s) + 4NO2(g) + O2(g) (M' = Li, Be, Mg, Ca, Sr, Ba) Thermal stability increases down the group (Be(NO3)2 < Mg(NO3)2 < Ca(NO3)2 < Sr(NO3)2 < Ba(NO3)2) as the polarizing power of the cation decreases with increasing size. ### (b) Carbonates (M2CO3 and M'CO3) **1. Solubility:** * **Alkali Metal Carbonates:** All alkali metal carbonates (e.g., Na2CO3, K2CO3) are soluble in water. Lithium carbonate (Li2CO3) is an exception, being sparingly soluble due to its high lattice energy (small Li+ and CO3^2-). * **Alkaline Earth Metal Carbonates:** Alkaline earth metal carbonates (e.g., CaCO3, MgCO3) are generally insoluble or sparingly soluble in water. This is due to their very high lattice energies, which are not sufficiently overcome by hydration energies. Solubility generally decreases down the group (MgCO3 > CaCO3 > SrCO3 > BaCO3) as lattice energy decreases less rapidly than hydration energy for large anions. **2. Thermal Stability:** * **Alkali Metal Carbonates:** Alkali metal carbonates are highly thermally stable. They decompose at very high temperatures, if at all. Lithium carbonate (Li2CO3) is an exception, being less stable and decomposing at a lower temperature to lithium oxide and carbon dioxide, similar to AEM carbonates: Li2CO3(s) → Li2O(s) + CO2(g) Other alkali metal carbonates are very stable and decompose only at extremely high temperatures (e.g., Na2CO3 melts without decomposition). * **Alkaline Earth Metal Carbonates:** Alkaline earth metal carbonates are much less thermally stable than alkali metal carbonates. The smaller, highly charged M2+ ions polarize the large carbonate ion (CO3^2-), weakening the C-O bonds and facilitating decomposition into metal oxide and carbon dioxide at relatively lower temperatures: M'CO3(s) → M'O(s) + CO2(g) Thermal stability increases down the group (BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3) as the polarizing power of the cation decreases with increasing size. ### (c) Sulphates (M2SO4 and M'SO4) **1. Solubility:** * **Alkali Metal Sulphates:** All alkali metal sulphates (e.g., Na2SO4, K2SO4) are highly soluble in water. Their hydration energies are sufficient to overcome their lattice energies. * **Alkaline Earth Metal Sulphates:** Alkaline earth metal sulphates show a decreasing trend in solubility down the group. MgSO4 is readily soluble, CaSO4 is sparingly soluble, and SrSO4 and BaSO4 are almost insoluble. This trend is explained by the fact that for large anions like SO4^2-, the lattice energy decreases less rapidly than the hydration energy as the cation size increases down the group. MgSO4 > CaSO4 > SrSO4 > BaSO4 (Solubility decreases) **2. Thermal Stability:** * **Alkali Metal Sulphates:** Alkali metal sulphates are very thermally stable and do not decompose easily even at high temperatures. * **Alkaline Earth Metal Sulphates:** Alkaline earth metal sulphates are generally thermally stable, but less so than their alkali metal counterparts. They decompose at high temperatures to form metal oxides, sulfur dioxide, and oxygen. Thermal stability generally increases down the group as the polarizing power of the cation decreases: M'SO4(s) → M'O(s) + SO2(g) + 1/2 O2(g) BeSO4 < MgSO4 < CaSO4 < SrSO4 < BaSO4 (Thermal stability increases)

Question 21

Q21. Describe the important uses of the following: (a) Caustic soda (b) Sodium carbonate (c) Quicklime.

Accepted Answer

The important uses of the given compounds are as follows:

### (a) Caustic Soda (NaOH)
Caustic soda, chemically known as sodium hydroxide, is a strong base with numerous industrial applications.
1.  **In the manufacture of soap and detergents:** It is used in the saponification process to convert fats and oils into soap.
2.  **In paper industry:** Used in the pulping process to dissolve lignin from wood fibers, producing pulp for paper manufacturing.
3.  **In textile industry:** Used for mercerizing cotton, which improves its strength, luster, and dye uptake.
4.  **In petroleum refining:** Used to remove acidic impurities (like H2S and mercaptans) from crude oil and petroleum products.
5.  **In the manufacture of artificial silk (rayon):** Used in the viscose process.
6.  **In the purification of bauxite:** Used in the Bayer process to extract alumina from bauxite ore.
7.  **As a laboratory reagent:** A common strong base used in various chemical reactions and titrations.

### (b) Sodium Carbonate (Na2CO3) - Washing Soda
Sodium carbonate, commonly known as washing soda, is an important industrial chemical.
1.  **In glass industry:** It is a primary raw material for the manufacture of glass, as it lowers the melting point of silica.
2.  **In soap and detergent industry:** Used as a builder in detergents to enhance their cleaning efficiency and as a filler in soap.
3.  **In paper industry:** Used in the pulping process and for bleaching paper.
4.  **For softening hard water:** It precipitates calcium and magnesium ions responsible for hardness, thus softening water.
5.  **In the manufacture of borax and caustic soda:** It is a starting material for the production of these chemicals.
6.  **As a laboratory reagent:** Used in various analytical and synthetic procedures.

### (c) Quicklime (CaO)
Quicklime, or calcium oxide, is a highly reactive basic oxide.
1.  **In cement industry:** It is a major component in the manufacture of cement, reacting with silica and alumina at high temperatures.
2.  **In steel industry:** Used as a flux in the basic oxygen steelmaking process to remove impurities like silica and phosphorus.
3.  **In agriculture:** Used to neutralize acidic soils (liming) and improve soil structure.
4.  **In sugar refining:** Used to purify sugar juice by precipitating impurities.
5.  **In the manufacture of calcium carbide (CaC2):** Used in electric arc furnaces to produce calcium carbide, which in turn is used to produce acetylene.
6.  **As a drying agent:** Its strong affinity for water makes it useful for drying gases and alcohols.
7.  **In the purification of water:** Used in some water treatment processes to adjust pH and remove impurities.

Question 22

Q22. Draw the structure of BeCl2 (vapour) and BeCl2 (solid).

Accepted Answer

### Structure of BeCl2 (Vapour Phase)
In the vapour phase, BeCl2 exists predominantly as a monomer at high temperatures and as a dimer at lower temperatures.

1.  **Monomeric BeCl2 (High Temperature):**
    *   **Geometry:** Linear
    *   **Hybridization:** sp hybridization of Be atom.
    *   **Bond Angle:** 180°
    *   **Description:** The beryllium atom forms two single covalent bonds with two chlorine atoms. Due to sp hybridization, the molecule has a linear geometry.

2.  **Dimeric BeCl2 (Lower Temperature):**
    *   **Geometry:** Bridged structure (non-planar)
    *   **Hybridization:** sp2 hybridization of Be atom.
    *   **Description:** Two BeCl2 units link together through chlorine bridges. Each beryllium atom is sp2 hybridized and is bonded to two terminal chlorine atoms and two bridging chlorine atoms. This forms a structure where two Be atoms are linked by two Cl atoms, forming a four-membered ring (Be-Cl-Be-Cl). The bridging chlorine atoms donate lone pairs to the electron-deficient beryllium atoms, forming coordinate bonds.

### Structure of BeCl2 (Solid Phase)
In the solid phase, BeCl2 exists as a polymeric chain structure.

*   **Geometry:** Polymeric chain
*   **Hybridization:** sp3 hybridization of Be atom.
*   **Description:** Each beryllium atom is tetrahedrally coordinated to four chlorine atoms. Two chlorine atoms act as terminal atoms, and two chlorine atoms act as bridging atoms, linking adjacent BeCl2 units. This forms an extended chain where each beryllium atom is surrounded by four chlorine atoms, and each chlorine atom bridges two beryllium atoms. The bridging bonds are dative bonds where chlorine donates a lone pair to the electron-deficient beryllium atom.

Question 23

Q23. Describe the important uses of the following: (a) Limestone (b) Cement (c) Plaster of Paris.

Accepted Answer

The important uses of the given compounds/materials are as follows:

### (a) Limestone (CaCO3)
Limestone is a sedimentary rock composed primarily of calcium carbonate (CaCO3). It is a versatile material with extensive applications.
1.  **In cement and concrete production:** It is a key raw material for the manufacture of Portland cement. When heated, it decomposes to quicklime (CaO), which is then used in cement production.
2.  **In steel industry:** Used as a flux in blast furnaces to remove impurities (like silica) from iron ore, forming slag.
3.  **In agriculture:** Used as an agricultural lime to neutralize acidic soils, improve soil structure, and provide calcium for plant growth.
4.  **As a building material:** Used as crushed stone for road construction, as dimension stone for building facades, and in the production of lime mortar.
5.  **In glass manufacturing:** Used as a source of calcium oxide, which improves the durability and strength of glass.
6.  **In desulfurization:** Used in flue gas desulfurization (FGD) systems to remove sulfur dioxide (SO2) from industrial exhaust gases.
7.  **In sugar refining:** Used to purify sugar by precipitating impurities.

### (b) Cement
Cement is a binder, a substance used for construction that sets, hardens, and adheres to other materials to bind them together. Portland cement is the most common type.
1.  **In concrete production:** Cement is the primary component of concrete, which is a mixture of cement, aggregates (sand and gravel), and water. Concrete is widely used for foundations, buildings, bridges, roads, and other infrastructure.
2.  **In mortar production:** Used to make mortar, a mixture of cement, sand, and water, which is used to bind bricks, blocks, and stones together in masonry construction.
3.  **In plaster and stucco:** Used in various plaster and stucco formulations for finishing walls and ceilings.
4.  **For grouting:** Used for filling gaps and cracks in structures, and for anchoring bolts and rebar.
5.  **In precast concrete products:** Used to manufacture precast concrete elements like pipes, blocks, panels, and railway sleepers.

### (c) Plaster of Paris (POP) (CaSO4.1/2H2O)
Plaster of Paris is calcium sulfate hemihydrate, obtained by heating gypsum (CaSO4.2H2O) to about 373 K. It has the unique property of setting into a hard mass upon mixing with water.
1.  **In construction and interior decoration:** Used for plastering walls and ceilings, creating decorative architectural elements, cornices, and false ceilings due to its smooth finish and quick setting properties.
2.  **In medical applications:** Used for setting fractured bones. When mixed with water, it forms a cast that immobilizes the fractured limb, aiding in healing.
3.  **In dentistry:** Used for making dental impressions and models of teeth.
4.  **For making molds and casts:** Used in sculpture, pottery, and for making molds for various objects due to its ability to capture fine details.
5.  **For making statues and decorative items:** Used to create intricate statues, figurines, and other ornamental pieces.
6.  **As a fire-proofing material:** Its hydrated form releases water upon heating, which helps in fire resistance.

Question 24

Q24. Explain the significance of the following: (a) Gypsum (b) Plaster of Paris (c) Limestone.

Accepted Answer

The significance of the following materials lies in their unique properties and extensive applications across various industries:

### (a) Gypsum (CaSO4.2H2O)
Gypsum, or calcium sulfate dihydrate, is a soft sulfate mineral. Its significance stems from its natural abundance and its role as a precursor to Plaster of Paris.
1.  **Source of Plaster of Paris:** Gypsum is the primary raw material for manufacturing Plaster of Paris (POP). When heated to about 373 K, it loses three-quarters of its water of crystallization to form POP.
    CaSO4.2H2O(s) → CaSO4.1/2H2O(s) + 3/2 H2O(g)
2.  **Retarder in Cement:** Gypsum is added to Portland cement during its manufacture. Its significance here is to slow down the setting time of cement, preventing 'flash setting' and allowing sufficient time for concrete to be mixed, transported, and placed before it hardens.
3.  **Fertilizer:** Used in agriculture as a soil conditioner and fertilizer, particularly for improving the structure of clay soils and providing calcium and sulfur, essential nutrients for plant growth.
4.  **Building Material:** Used in the manufacture of gypsum boards (drywall), which are widely used for walls and ceilings in construction due to their fire-resistant and sound-insulating properties.
5.  **Art and Sculpture:** Historically used for making molds and casts.

### (b) Plaster of Paris (POP) (CaSO4.1/2H2O)
Plaster of Paris, or calcium sulfate hemihydrate, is significant due to its unique property of setting into a hard, solid mass upon mixing with water, making it highly versatile.
1.  **Medical Applications:** Its most significant use is in orthopedics for setting fractured bones. When mixed with water, it forms a paste that can be molded around a fractured limb. Upon setting, it forms a rigid cast that immobilizes the bone, facilitating healing.
2.  **Construction and Decoration:** Widely used in building construction for plastering walls and ceilings, creating smooth surfaces, and for decorative purposes like cornices, false ceilings, and ornamental designs. Its quick setting and fine finish are highly valued.
3.  **Dental Impressions:** Crucial in dentistry for making accurate impressions of teeth and gums, which are then used to create dental models, crowns, and bridges.
4.  **Molds and Casts:** Its ability to capture intricate details makes it invaluable in art and sculpture for creating molds and casts of statues, figurines, and other decorative items.
5.  **Fire-proofing Material:** Due to its hydrated nature, it acts as a fire retardant. When exposed to heat, it releases water, which helps to cool the surface and inhibit combustion.

### (c) Limestone (CaCO3)
Limestone, primarily calcium carbonate, is one of the most significant and abundant natural resources. Its importance spans construction, industry, agriculture, and environmental applications.
1.  **Foundation of Construction Industry:** It is the primary raw material for the production of quicklime (CaO) and Portland cement, which are fundamental to concrete and mortar. Thus, it underpins modern infrastructure development.
2.  **Metallurgical Flux:** In the steel industry, it is indispensable as a flux in blast furnaces. It reacts with impurities like silica and phosphorus to form slag, purifying the molten metal.
3.  **Soil Amendment:** In agriculture, it is used to neutralize acidic soils (liming), improving soil fertility and crop yields. It also supplies essential calcium to plants.
4.  **Environmental Protection:** Significant in flue gas desulfurization (FGD) systems, where it reacts with sulfur dioxide (SO2) from industrial emissions, reducing air pollution.
5.  **Glass Manufacturing:** An essential component in glass production, where it contributes to the stability and durability of the glass.
6.  **Chemical Industry Raw Material:** Used in the production of various chemicals, including soda ash (sodium carbonate) and calcium carbide.
7.  **Building and Decorative Stone:** Used as crushed aggregate for roads and as dimension stone for architectural purposes due to its aesthetic appeal and durability.

Question 25

Q25. Write a balanced chemical equation for each of the following reactions: (a) Burning of sodium in air (b) Heating of sodium nitrate (c) Reaction of potassium with water (d) Reaction

Accepted Answer

Here are the balanced chemical equations for the given reactions:

(a) **Burning of sodium in air:**
When sodium burns in air (which contains both O2 and N2), it primarily forms sodium peroxide (Na2O2) and also reacts with nitrogen to form sodium nitride (Na3N).

*   **With Oxygen:**
    2Na(s) + O2(g) → Na2O2(s)

*   **With Nitrogen (minor reaction):**
    6Na(s) + N2(g) → 2Na3N(s)

(b) **Heating of sodium nitrate:**
When sodium nitrate is heated, it decomposes to form sodium nitrite and oxygen gas.

2NaNO3(s) → 2NaNO2(s) + O2(g)

(c) **Reaction of potassium with water:**
Potassium reacts vigorously and exothermically with water to produce potassium hydroxide and hydrogen gas. The reaction is so exothermic that the hydrogen gas produced often ignites.

2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)

(d) **Reaction of steam with magnesium:**
Magnesium reacts with steam (gaseous water) at high temperatures to form magnesium oxide and hydrogen gas.

Mg(s) + H2O(g) → MgO(s) + H2(g)

(e) **Burning of calcium in air:**
When calcium burns in air, it primarily forms calcium oxide (CaO) and also reacts with nitrogen to form calcium nitride (Ca3N2).

*   **With Oxygen:**
    2Ca(s) + O2(g) → 2CaO(s)

*   **With Nitrogen (minor reaction):**
    3Ca(s) + N2(g) → Ca3N2(s)

Question 26

Q26. Describe the important uses of the following: (a) Caustic soda (b) Sodium carbonate (c) Quicklime.

Accepted Answer

The important uses of the given compounds are as follows:

**(a) Caustic soda (NaOH)**
Caustic soda, chemically known as sodium hydroxide, is a highly versatile compound with numerous industrial applications due to its strong basic nature.
*   **Soap and Detergent Industry:** It is a key ingredient in the saponification process for making soaps and detergents. It reacts with fats and oils to produce soap.
*   **Paper Industry:** Used in the pulping process to dissolve lignin from wood, separating cellulose fibers for paper production.
*   **Textile Industry:** Employed in mercerizing cotton to improve its strength, luster, and dye uptake.
*   **Petroleum Refining:** Used to remove acidic impurities (like H2S and mercaptans) from crude oil and petroleum products.
*   **Alumina Production:** Essential in the Bayer process for refining bauxite ore to produce alumina (Al2O3), which is then used to make aluminium metal.
*   **Water Treatment:** Used for pH adjustment and to precipitate heavy metals in wastewater treatment.
*   **Chemical Manufacturing:** A fundamental raw material in the production of various chemicals, including dyes, pharmaceuticals, and other sodium compounds.

**(b) Sodium carbonate (Na2CO3)**
Sodium carbonate, also known as soda ash, is another industrially significant compound.
*   **Glass Industry:** It is a primary raw material in the manufacture of glass, where it lowers the melting point of silica, reducing energy consumption.
*   **Soap and Detergent Industry:** Used as a builder in detergents to enhance their cleaning efficiency by softening water and emulsifying grease.
*   **Paper Industry:** Employed in the pulping process and for bleaching paper.
*   **Water Treatment:** Used for softening hard water by precipitating calcium and magnesium ions as carbonates.
*   **Textile Industry:** Used in dyeing and printing processes.
*   **Chemical Manufacturing:** A source of sodium ions for the production of various sodium compounds like sodium silicate, sodium bicarbonate, and sodium chromate.
*   **Metallurgy:** Used as a flux in some metallurgical processes.

**(c) Quicklime (CaO)**
Quicklime, or calcium oxide, is a widely used alkaline compound with applications primarily in construction and chemical industries.
*   **Cement and Mortar Production:** It is a crucial component in the manufacture of cement (Portland cement) and mortar. It reacts with water to form slaked lime, which is then used in construction.
*   **Steel Manufacturing:** Used as a flux in the basic oxygen steelmaking process to remove impurities like silica, phosphorus, and sulfur from molten iron.
*   **Agriculture:** Applied to acidic soils as a liming agent to neutralize acidity, improve soil structure, and provide calcium nutrients.
*   **Water Treatment:** Used for softening hard water (Clark's process) and for pH adjustment in wastewater treatment.
*   **Chemical Industry:** A raw material for producing calcium carbide, bleaching powder, and other calcium compounds.
*   **Sugar Refining:** Used to clarify sugar beet and sugarcane juices by precipitating impurities.
*   **Glass Industry:** Used in the manufacture of some types of glass.

Question 27

Q27. When a metal X is treated with sodium hydroxide, a white precipitate (A) is obtained, which is soluble in excess of sodium hydroxide to give soluble complex (B). Compound (A) is so

Accepted Answer

Let's identify the compounds and write the relevant equations step-by-step.

**Identification:**
The key information is that metal X forms a white precipitate (A) with NaOH, which is soluble in excess NaOH to form a soluble complex (B), and also soluble in dilute HCl to form compound (C). This indicates that compound (A) is an **amphoteric hydroxide**. Among the s-block elements, **Beryllium (Be)** forms an amphoteric hydroxide, Be(OH)2.

Let's verify the properties with Beryllium:
*   **Metal X:** Beryllium (Be)
*   **Reaction with NaOH to form precipitate (A):** Beryllium salts (e.g., BeCl2) react with NaOH to form beryllium hydroxide precipitate.
    BeCl2(aq) + 2NaOH(aq) → Be(OH)2(s) + 2NaCl(aq)
    So, **A = Be(OH)2 (Beryllium hydroxide)**

*   **Solubility of (A) in excess NaOH to give soluble complex (B):** Beryllium hydroxide is amphoteric and reacts with excess strong base.
    Be(OH)2(s) + 2NaOH(aq) (excess) → Na2[Be(OH)4](aq)
    So, **B = Na2[Be(OH)4] (Sodium tetrahydroxoberyllate(II))**

*   **Solubility of (A) in dilute HCl to form compound (C):** Beryllium hydroxide is amphoteric and reacts with acid.
    Be(OH)2(s) + 2HCl(aq) → BeCl2(aq) + 2H2O(l)
    So, **C = BeCl2 (Beryllium chloride)**

*   **Heating (A) strongly to give oxide (D) used in chromatography:** Beryllium hydroxide decomposes on strong heating to form beryllium oxide.
    Be(OH)2(s) --(heat)--> BeO(s) + H2O(g)
    So, **D = BeO (Beryllium oxide)**

*Note on chromatography use:* While aluminium oxide (Al2O3) is a very common adsorbent in chromatography, beryllium oxide (BeO) also possesses properties (e.g., high surface area, thermal stability) that make it suitable for specialized chromatographic applications, especially in gas chromatography or as a support material. Given that the question is from the s-block chapter, Beryllium is the most appropriate s-block metal that fits all the described chemical properties.

**Summary of Identifications:**
*   **X = Beryllium (Be)**
*   **A = Beryllium hydroxide (Be(OH)2)**
*   **B = Sodium tetrahydroxoberyllate(II) (Na2[Be(OH)4])**
*   **C = Beryllium chloride (BeCl2)**
*   **D = Beryllium oxide (BeO)**

**Relevant Equations:**
1.  **Formation of precipitate (A):**
    BeCl2(aq) + 2NaOH(aq) → Be(OH)2(s) + 2NaCl(aq)
    *(Assuming a soluble beryllium salt reacts with NaOH to form the precipitate)*

2.  **Solubility of (A) in excess NaOH to form (B):**
    Be(OH)2(s) + 2NaOH(aq) (excess) → Na2[Be(OH)4](aq)

3.  **Solubility of (A) in dilute HCl to form (C):**
    Be(OH)2(s) + 2HCl(aq) → BeCl2(aq) + 2H2O(l)

4.  **Decomposition of (A) to form (D):**
    Be(OH)2(s) --(heat)--> BeO(s) + H2O(g)

Question 28

Q28. Describe the structure of the following: (a) Beryllium chloride (vapour) (b) Beryllium chloride (solid).

Accepted Answer

Beryllium chloride (BeCl2) exhibits different structures in its vapour and solid phases due to the varying conditions and the tendency of beryllium to achieve a stable electron configuration.

**(a) Beryllium chloride (vapour phase)**
In the vapour phase, the structure of beryllium chloride depends on the temperature.
*   **At high temperatures (above 1200 K):** Beryllium chloride exists as discrete **monomeric** BeCl2 units. In this monomeric form, beryllium is sp hybridized, and the molecule has a **linear geometry** with a Cl-Be-Cl bond angle of 180°. Beryllium has only four electrons in its valence shell, making it electron-deficient.
*   **At lower temperatures (around 1000 K):** The monomeric units associate to form a **dimeric** structure, (BeCl2)2. In the dimer, two BeCl2 units are linked by two chlorine atoms forming **chlorine bridges**. Each beryllium atom is sp2 hybridized and is surrounded by three chlorine atoms (two terminal and one bridging), resulting in a **trigonal planar** arrangement around each beryllium. The bridging chlorine atoms donate lone pairs to the electron-deficient beryllium atoms, forming coordinate covalent bonds.

**(b) Beryllium chloride (solid phase)**
In the solid state, beryllium chloride forms a **polymeric chain structure**. This structure is achieved by each beryllium atom being covalently bonded to four chlorine atoms, forming a tetrahedral arrangement around each beryllium. Each chlorine atom acts as a bridge between two beryllium atoms.
*   Each beryllium atom is **sp3 hybridized**.
*   The structure consists of an infinite chain where each beryllium atom is surrounded by four chlorine atoms in a **tetrahedral geometry**.
*   Each chlorine atom forms two single bonds, acting as a **bridging ligand** between two beryllium atoms.
*   This bridging allows the electron-deficient beryllium atoms to complete their octet (or achieve a stable electron count of 8 electrons in their valence shell through coordinate bonds from chlorine lone pairs), increasing the stability of the compound.
*   The repeating unit of the chain can be represented as (-BeCl2-)n.

These structural variations highlight the electron-deficient nature of beryllium and its tendency to achieve coordination saturation through polymerization and bridging.

Question 29

Q29. What is the significance of the diagonal relationship between Li and Mg, and Be and Al?

Accepted Answer

The diagonal relationship is a phenomenon observed in the periodic table where elements of the second period (Li, Be, B) show similarities in properties with elements of the third period in the next group (Mg, Al, Si, respectively). This similarity arises because, on moving diagonally across the periodic table, the effects of decreasing atomic size and increasing nuclear charge (across a period) are somewhat compensated by the effect of increasing atomic size and decreasing nuclear charge (down a group). This results in comparable electronegativity, polarizing power, and charge-to-radius ratio (ionic potential) for diagonally related elements.

**Significance of Diagonal Relationship between Li and Mg:**
Lithium (Li) from Group 1 and Magnesium (Mg) from Group 2 exhibit significant similarities, which are distinct from other alkali metals and alkaline earth metals, respectively.

*   **Similarities in Properties:**
    1.  **Atomic and Ionic Sizes:** Li+ (76 pm) and Mg2+ (72 pm) have very similar ionic radii.
    2.  **Electronegativity:** Both have similar electronegativity values (Li: 0.98, Mg: 1.31).
    3.  **Polarizing Power:** Both Li+ and Mg2+ ions have high and comparable polarizing power due to their small size and relatively high charge (Li+ has higher charge density than other alkali metals, Mg2+ has higher charge density than other alkaline earth metals). This leads to a greater degree of covalent character in their compounds.
    4.  **Hardness:** Both Li and Mg are harder and lighter than other elements in their respective groups.
    5.  **Reactivity:** Both react slowly with water (compared to other alkali metals for Li, and other alkaline earth metals for Mg).
    6.  **Oxides and Hydroxides:** Both form normal oxides (Li2O, MgO) and their hydroxides (LiOH, Mg(OH)2) are weak bases and sparingly soluble in water. They decompose on heating.
    7.  **Carbonates:** Both Li2CO3 and MgCO3 are thermally unstable and decompose on heating to form oxides and CO2 (Li2CO3 → Li2O + CO2; MgCO3 → MgO + CO2). Other alkali metal carbonates are stable.
    8.  **Nitrates:** Both LiNO3 and Mg(NO3)2 decompose on heating to give oxides, NO2, and O2 (LiNO3 → Li2O + NO2 + O2; Mg(NO3)2 → MgO + NO2 + O2). Other alkali metal nitrates give nitrites.
    9.  **Chlorides:** Both LiCl and MgCl2 are deliquescent (absorb moisture) and crystallize from aqueous solutions as hydrates (LiCl·2H2O, MgCl2·6H2O). They are soluble in organic solvents like ethanol.
    10. **Formation of Nitrides:** Both Li and Mg react directly with nitrogen to form nitrides (Li3N, Mg3N2). Other alkali metals do not form nitrides directly.

**Significance of Diagonal Relationship between Be and Al:**
Beryllium (Be) from Group 2 and Aluminium (Al) from Group 13 also exhibit striking similarities.

*   **Similarities in Properties:**
    1.  **Atomic and Ionic Sizes:** Be2+ (31 pm) and Al3+ (53.5 pm) have similar charge-to-radius ratios, leading to similar polarizing power.
    2.  **Electronegativity:** Both have similar electronegativity values (Be: 1.57, Al: 1.61).
    3.  **Covalent Character:** Both form compounds with significant covalent character due to high polarizing power of Be2+ and Al3+ ions.
    4.  **Amphoteric Nature:** Both Be and Al form amphoteric oxides (BeO, Al2O3) and hydroxides (Be(OH)2, Al(OH)3), meaning they react with both acids and strong bases.
    5.  **Reactivity with Acids/Bases:** Both metals dissolve in strong acids and strong bases, liberating hydrogen gas.
        Be + 2NaOH + 2H2O → Na2[Be(OH)4] + H2
        2Al + 2NaOH + 6H2O → 2Na[Al(OH)4] + 3H2
    6.  **Passivity:** Both metals are rendered passive by nitric acid due to the formation of a protective oxide layer on their surface.
    7.  **Chlorides:** Both BeCl2 and AlCl3 are Lewis acids and exist as bridged polymers in the solid state and as dimers in the vapour phase. They are covalent and soluble in organic solvents.
    8.  **Complex Formation:** Both Be2+ and Al3+ ions have a strong tendency to form complex compounds.

In essence, the diagonal relationship helps in understanding why certain elements in the second period deviate from the typical group trends and instead resemble elements in the next group of the third period, providing a deeper insight into periodic properties and chemical behavior.

Question 30

Q30. Why are lithium salts commonly hydrated and other alkali metal ions are mostly anhydrous?

Accepted Answer

Lithium salts are commonly hydrated, while other alkali metal ions are mostly anhydrous due to the exceptionally small size and high charge density of the Li⁺ ion.

1.  **Small Ionic Size:** Among all alkali metal ions, Li⁺ has the smallest ionic radius (approx. 76 pm).
2.  **High Charge Density:** Due to its small size, the positive charge of the Li⁺ ion is concentrated over a very small volume, leading to a very high charge density (charge/radius ratio).
3.  **Strong Electrostatic Attraction:** This high charge density enables the Li⁺ ion to exert a strong electrostatic attraction on the lone pair electrons of the oxygen atom in water molecules.
4.  **High Hydration Enthalpy:** Consequently, Li⁺ ions have a very high hydration enthalpy (energy released when ions are hydrated). This high hydration enthalpy is sufficient to overcome the lattice energy of many lithium salts and form stable hydrated compounds, such as LiCl·2H₂O.

In contrast, other alkali metal ions (Na⁺, K⁺, Rb⁺, Cs⁺) are much larger in size. As their size increases down the group, their charge density decreases significantly. This results in weaker electrostatic attraction towards water molecules and lower hydration enthalpies, making them less prone to forming hydrated salts; hence, they are mostly anhydrous.

Question 31

Q31. Why is LiF almost insoluble in water whereas LiCl is soluble not only in water but also in acetone?

Accepted Answer

The difference in solubility between LiF and LiCl can be explained by considering their lattice energy, hydration energy, and covalent character.

**1. LiF is almost insoluble in water:**
*   **High Lattice Energy:** LiF has an exceptionally high lattice energy. This is because both Li⁺ and F⁻ ions are very small in size. The strong electrostatic forces of attraction between these small, oppositely charged ions result in a very stable crystal lattice that requires a large amount of energy to break.
*   **Insufficient Hydration Energy:** Although Li⁺ has a high hydration energy, the hydration energy of both Li⁺ and F⁻ ions combined is not sufficient to overcome the extremely high lattice energy of LiF. Therefore, LiF does not readily dissolve in water.

**2. LiCl is soluble in water:**
*   **Lower Lattice Energy:** Compared to F⁻, the Cl⁻ ion is much larger. This larger size of Cl⁻ leads to weaker electrostatic interactions with Li⁺, resulting in a significantly lower lattice energy for LiCl compared to LiF.
*   **Sufficient Hydration Energy:** The hydration energy of Li⁺ and Cl⁻ ions is sufficient to overcome the relatively lower lattice energy of LiCl. This makes LiCl readily soluble in water.

**3. LiCl is soluble in acetone:**
*   **Covalent Character:** According to Fajan's rules, a small cation (Li⁺) and a relatively large anion (Cl⁻) lead to significant covalent character in the bond. The Li⁺ ion can polarize the electron cloud of the larger Cl⁻ ion more effectively, imparting partial covalent character to the Li-Cl bond.
*   **Solubility in Organic Solvents:** Due to this appreciable covalent character, LiCl is not purely ionic. It can dissolve not only in highly polar solvents like water but also in moderately polar organic solvents like acetone, which can solvate the partially covalent LiCl molecules.

Question 32

Q32. Explain the significance of the following:
(a) Gypsum
(b) Plaster of Paris
(c) Milk of magnesia.

Accepted Answer

The compounds listed have significant applications in various fields due to their unique chemical and physical properties.

### (a) Gypsum (CaSO₄·2H₂O)
Gypsum is a naturally occurring mineral, chemically known as calcium sulfate dihydrate. Its significance stems from its widespread availability and diverse industrial applications.

**Significance and Uses:**
1.  **Raw Material for Plaster of Paris (POP):** Gypsum is the primary raw material for the manufacture of Plaster of Paris. When heated to about 373 K (100 °C), it loses three-quarters of its water of crystallization to form POP.
    `CaSO₄·2H₂O (Gypsum) --(Heat to 373 K)--> CaSO₄·½H₂O (POP) + 1½H₂O`
2.  **Cement Industry:** It is added to cement during its manufacture. Its main role is to act as a retarder, slowing down the setting time of cement. This allows for proper mixing, transportation, and application of the concrete before it hardens.
3.  **Agriculture:** Gypsum is used as a soil conditioner, particularly in alkaline or sodic soils. It improves soil structure, water penetration, and provides essential calcium and sulfur nutrients to plants.
4.  **Building Materials:** It is used in the production of drywall, plasterboards, and as a component in some building blocks and ceiling tiles due to its fire-resistant properties and ability to provide a smooth finish.

### (b) Plaster of Paris (POP) (CaSO₄·½H₂O)
Plaster of Paris, or calcium sulfate hemihydrate, is a white powdery material known for its unique property of setting into a hard mass upon mixing with water. This property makes it invaluable in several applications.

**Significance and Uses:**
1.  **Medical Field:** It is extensively used in medicine for setting fractured bones. When mixed with water, it forms a paste that can be molded around the fractured limb. Upon drying, it hardens into a rigid cast, immobilizing the bone and aiding in healing. It is also used in dentistry for making molds and impressions.
2.  **Construction and Decoration:** POP is widely used in construction for plastering walls and ceilings to give a smooth, even finish. It is also used for making decorative articles, statues, false ceilings, and ornamental designs due to its ability to take intricate shapes.
3.  **Fire-proofing Material:** Due to its non-combustible nature and the release of water vapor when heated, POP acts as a good fire-resistant material, often used in fire-proofing applications.
4.  **Laboratory Uses:** It is used in laboratories for making casts, molds, and for sealing air-tight apparatus.

### (c) Milk of Magnesia (Mg(OH)₂ suspension)
Milk of Magnesia is a common name for a suspension of magnesium hydroxide in water. It is a mild base and its significance lies primarily in its medicinal properties.

**Significance and Uses:**
1.  **Antacid:** Its most significant use is as an antacid. Magnesium hydroxide is a weak base that neutralizes excess hydrochloric acid (HCl) in the stomach, providing relief from heartburn, indigestion, and sour stomach. It does not cause gas formation, unlike some other antacids.
    `Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)`
2.  **Laxative:** In higher doses, Milk of Magnesia acts as a mild osmotic laxative. It draws water into the intestines, softening the stool and promoting bowel movements, thereby relieving constipation.
3.  **Topical Use:** Occasionally, it is used topically to relieve minor skin irritations or rashes due to its soothing and drying properties.

Chapter 10: The s-Block Elements