P-Block Elements (LAQs)

Chemistry-2 | 6. P-Block Elements – LAQs:
Welcome to LAQs in Chapter 6: P-Block Elements. This page features the important FAQs for Long Answer Questions. Each answer is provided in simple English, with a Telugu explanation, and formatted according to the exam style. This will help you thoroughly prepare and secure top marks in your final exams.

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LAQ-1 : How is ammonia manufactured by Haber’s process?

Introduction

The Haber’s process is a cornerstone industrial method for synthesizing ammonia (NH₃) from nitrogen (N₂) and hydrogen (H₂) gases, employing direct synthesis under specific conditions.

Chemical Reaction

1. The formation of ammonia is represented by the equation: $$N_2 (g) + 3H_2 (g) \rightleftharpoons 2NH_3 (g)$$
2. This reaction is exothermic, with a change in enthalpy (ΔH) of -92.3 kJ.

Factors Affecting the Equilibrium

1. Effect of Temperature: Lower temperatures favor the forward reaction, but an intermediate temperature is chosen to balance reaction rate and yield.
2. Effect of Pressure: Higher pressure shifts the equilibrium towards ammonia production due to a decrease in volume.

Optimum Conditions for Production

1. Temperature: 725-775 K (approx. 450-500°C)
2. Pressure: 200-300 atm
3. Catalyst: Finely divided iron (Fe), enhanced by a promoter, molybdenum (Mo).

Process Steps

1. Compression of Gases: Nitrogen and hydrogen gases are compressed to 200-300 atm.
2. Passing Through Catalyst Chamber: The compressed gases pass over the iron catalyst at 500°C, with molybdenum serving as the promoter.
3. Cooling and Liquefaction: Resultant gases are cooled, causing ammonia to liquefy and be collected.
4. Recirculation of Unreacted Gases: Unreacted nitrogen and hydrogen are recycled back into the system for efficiency.

Summary

The Haber’s process ingeniously applies the principles of chemical equilibrium to optimize ammonia production conditions, demonstrating a significant advancement in industrial chemistry. This method ensures efficient, large-scale production of ammonia, pivotal for various applications, including fertilizers and chemicals.

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LAQ-2 : How does ammonia react with i) ZnSO₄(aq) ii) CuSO₄(aq) iii) AgCl_(s)

Introduction

Ammonia (NH₃) is a versatile compound that reacts differently with various substances. Below are the reactions of ammonia with zinc sulfate (ZnSO₄), copper sulfate (CuSO₄), and silver chloride (AgCl).

1. Reaction with Zinc Sulfate (ZnSO₄)
   • Reaction: Ammonia reacts with ZnSO₄(aq) to form a complex compound. $$ZnSO_4(aq) + 4NH_3(aq) \rightarrow [Zn(NH_3)_4]SO_4(aq)$$
   • Result: The formation of tetraamminezinc(II) sulfate is observed, indicating the complexation of zinc ions with ammonia.
2. Reaction with Copper Sulfate (CuSO₄)
   • Reaction: Ammonia interacts with CuSO₄(aq), leading to the formation of a deep blue complex. $$CuSO_4(aq) + 4NH_3(aq) \rightarrow [Cu(NH_3)_4]SO_4(aq)$$
   • Result: The reaction produces tetraamminecopper(II) sulfate, recognized by its characteristic deep blue color, which indicates the complexation of copper ions with ammonia.
3. Reaction with Silver Chloride (AgCl)
   • Reaction: Ammonia dissolves AgCl(s), forming a soluble complex. $$AgCl(s) + 2NH_3(aq) \rightarrow [Ag(NH_3)_2]Cl(aq)$$
   • Result: The formation of diamminesilver(I) chloride demonstrates the ability of ammonia to dissolve otherwise insoluble silver chloride by complexation.

Summary

Ammonia’s reactions with zinc sulfate, copper sulfate, and silver chloride showcase its reactive nature and ability to form complex compounds. These reactions highlight the importance of ammonia in analytical chemistry, particularly in the formation of complex ions with metals such as zinc, copper, and silver.

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LAQ-3 : How is nitric acid manufactured by Ostwald’s process?

Introduction

The Ostwald’s Process is a principal method for producing nitric acid by catalytically oxidizing ammonia. This process is efficient and widely used in the chemical industry.

Stepwise Procedure

1. Oxidation of Ammonia to Nitric Oxide
   • Ammonia is oxidized to nitric oxide using a platinum catalyst at approximately 500°C. $$4NH_3(g) + 5O_2(g) \rightarrow 4NO(g) + 6H_2O(g)$$
   • The platinum gauze acts as a catalyst, enhancing the reaction rate.
2. Oxidation of Nitric Oxide to Nitrogen Dioxide: The formed nitric oxide is further oxidized to nitrogen dioxide in the presence of oxygen. $$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$$
3. Absorption of Nitrogen Dioxide to form Nitric Acid
   • Nitrogen dioxide is cooled and then absorbed in water, producing nitric acid. $$3NO_2(g) + H_2O(l) \rightarrow 2HNO_3(aq) + NO(g)$$
   • A portion of nitric oxide is recycled, enhancing process efficiency.

Concentration of Nitric Acid

1. First Stage: Dilute nitric acid is distilled to obtain about 68% HNO₃ by mass.
2. Second Stage: The 68% nitric acid is further dehydrated with concentrated sulfuric acid, yielding anhydrous HNO₃ with a concentration of around 98%.
3. Third Stage: To achieve 100% nitric acid, the 98% acid is cooled in a freezing mixture, leading to the crystallization of pure HNO₃.

Summary

The Ostwald process efficiently produces nitric acid, a colorless, strong oxidizing liquid used across various industries and laboratories. Through the catalytic oxidation of ammonia and subsequent processing stages, high-purity nitric acid is obtained, demonstrating the importance of this method in the chemical manufacturing sector.

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LAQ-4 : How does nitric acid react with the following? i) Copper ii) Zn iii) S₈ iv) P₄

Introduction

Nitric acid (HNO₃) is a highly reactive, strong acid that can undergo various chemical reactions with metals, non-metals, and other compounds.

1. Reaction with Copper (Cu)
   • Reaction: Nitric acid reacts with copper to produce copper(II) nitrate, nitrogen dioxide, and water. $$3Cu + 8HNO_3 \rightarrow 3Cu(NO_3)_2 + 2NO + 4H_2O$$
   • Copper is oxidized to copper(II) nitrate, while nitric acid is reduced to nitrogen dioxide (NO) gas.
2. Reaction with Zinc (Zn)
   • Reaction: When zinc reacts with nitric acid, zinc nitrate, nitrogen monoxide, and water are produced. $$Zn + 4HNO_3 \rightarrow Zn(NO_3)_2 + 2NO_2 + 2H_2O$$
   • Zinc is oxidized to zinc nitrate, and nitric acid is reduced to nitrogen dioxide.
3. Reaction with Sulfur (S₈)
   • Reaction: Nitric acid oxidizes sulfur to sulfuric acid, producing nitrogen dioxide as a byproduct. $$S_8 + 48HNO_3 \rightarrow 8H_2SO_4 + 48NO_2 + 16H_2O$$
   • Sulfur is oxidized to sulfuric acid, and nitric acid is reduced to nitrogen dioxide.
4. Reaction with Phosphorus (P₄)
   • Reaction: Phosphorus reacts with nitric acid to form phosphoric acid, nitrogen dioxide, and water. $$P_4 + 20HNO_3 \rightarrow 4H_3PO_4 + 20NO_2 + 4H_2O$$
   • Phosphorus is oxidized to phosphoric acid, and nitric acid is reduced to nitrogen dioxide.

Summary

Nitric acid’s reactions with copper, zinc, sulfur, and phosphorus illustrate its strong oxidizing properties. These reactions produce various compounds, such as nitrates and oxides, demonstrating nitric acid’s versatility in chemical transformations.

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LAQ-5 : a. How is Ozone prepared?
b. How does it react with i) C₂H₄ ii) KL iii) Hg iv) Pbs v) NO vi) Ag

Introduction

Ozone (O₃) is a powerful oxidizing agent with significant applications in both natural and industrial processes.

a. Preparation of Ozone

1. Ozone is primarily prepared by passing an electric discharge through oxygen gas (O₂). This process, known as silent electric discharge, converts oxygen molecules into ozone. $$3O_2 \overset{electric\ discharge}{\longrightarrow} 2O_3$$
2. The formation of ozone from oxygen involves the breakdown of O₂ molecules and the recombination of oxygen atoms to form O₃.

b. Reactions of Ozone

1. With Ethylene (C₂H₄)
   • Reaction: Ozone oxidizes ethylene to form ethyl ozonide, which can further decompose to form aldehydes or ketones. $$C_2H_4 + O_3 \rightarrow C_2H_4O_3$$
2. With Potassium Iodide (KI)
   • Reaction: Ozone reacts with potassium iodide in the presence of water, releasing iodine, which can be detected by its brown color. $$2KI + O_3 + H_2O \rightarrow I_2 + 2KOH + O_2$$
3. With Mercury (Hg)
   • Reaction: Ozone oxidizes mercury to mercuric oxide (HgO), demonstrating its strong oxidizing ability. $$Hg + O_3 \rightarrow HgO + O_2$$
4. With Lead Sulfide (PbS)
   • Reaction: Ozone oxidizes lead sulfide to lead sulfate. $$PbS + 4O_3 \rightarrow PbSO_4 + 4O_2$$
5. With Nitric Oxide (NO)
   • Reaction: Ozone readily oxidizes nitric oxide to nitrogen dioxide (NO₂), a process used in air purification. $$NO + O_3 \rightarrow NO_2 + O_2$$
6. With Silver (Ag)
   • Reaction: Ozone reacts with silver to form silver oxide (Ag₂O), showcasing its ability to oxidize metals. $$2Ag + O_3 \rightarrow Ag_2O + O_2$$

Summary

The preparation and reactions of ozone highlight its potent oxidizing nature. Whether transforming organic compounds, reacting with metals, or purifying air by oxidizing pollutants, ozone’s chemical behavior underscores its importance across various fields.

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LAQ-6 : Describe the manufacture of H₂SO₄ by contact process.

Introduction

The Contact Process is a prevalent and efficient method for producing sulphuric acid (H₂SO₄), essential for various industrial applications. This multi-step process involves the oxidation of sulphur dioxide to sulphur trioxide, followed by its conversion to sulphuric acid.

Steps in the Contact Process

1. Formation of Sulphur Dioxide (SO₂)
   • Burning Sulphur: Sulphur is combusted in air, yielding sulphur dioxide. $$\textbf{S} + \textbf{O}_2 \rightarrow \textbf{SO}_2​$$
2. Conversion of SO₂ to Sulphur Trioxide (SO₃)
   • Oxidizing SO₂: Sulphur dioxide is oxidized to sulphur trioxide, an exothermic reaction. $$2\textbf{SO}_2(g) + \textbf{O}_2(g) \rightarrow 2\textbf{SO}_3(g)$$
   • Optimal Conditions: To favor SO₃ formation, conditions include a temperature of 673K to 723K, pressure of 1.5 to 2 atm, and vanadium(V) oxide (V₂O₅) or platinum (Pt) as catalysts.
3. Formation of Oleum
   • Absorbing SO₃: Sulphur trioxide is absorbed by 98% concentrated sulphuric acid to form oleum. $$\textbf{SO}_3 + \textbf{H}_2\textbf{SO}_4 \rightarrow \textbf{H}_2\textbf{S}_2\textbf{O}_7$$
4. Production of Concentrated Sulphuric Acid
   • Diluting Oleum: Oleum is then diluted with water to yield concentrated sulphuric acid. $$\textbf{H}_2\textbf{S}_2\textbf{O}_7 + \textbf{H}_2\textbf{O} \rightarrow 2\textbf{H}_2\textbf{SO}_4$$

Summary

The Contact Process effectively produces sulphuric acid through the combustion of sulphur, oxidation of SO₂ to SO₃, and subsequent absorption of SO₃ in sulphuric acid to form oleum. The careful management of reaction conditions ensures high yields of sulphuric acid, highlighting the process’s significance in the chemical industry.

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LAQ-7 : How is chlorine prepared in the laboratory?
How chlorine is prepared in Deacon’s process
How does it react with the following.
i) Cold dil. NaOH ii) Hot conc. NaOH iii) Excess NH₃ iv) Slaked lime
v) Hypo (Na₂ S₂ O₃)

Introduction

Chlorine is a highly reactive, greenish-yellow gas, widely used for water purification, in the production of various chemicals, and as a bleaching agent.

1. Preparation of Chlorine in the Laboratory
   • Laboratory Method: Chlorine can be prepared by the reaction of manganese dioxide (MnO₂) with hydrochloric acid (HCl). $$\textbf{MnO}_2 + 4\textbf{HCl} \rightarrow \textbf{MnCl}_2 + 2\textbf{H}_2\textbf{O} + \textbf{Cl}_2$$
   • Manganese dioxide acts as a catalyst, facilitating the liberation of chlorine gas.
2. Preparation of Chlorine by Deacon’s Process
   • Deacon’s Process: Involves the catalytic oxidation of hydrochloric acid with oxygen in the presence of a copper(II) chloride (CuCl₂) catalyst at about 450°C. $$4\textbf{HCl} + \textbf{O}_2 \overset{\textbf{CuCl}_2}{\longrightarrow} 2\textbf{Cl}_2 + 2\textbf{H}_2\textbf{O}$$

Reactions of Chlorine

1. With Cold Dilute NaOH
   • Reaction: Forms sodium hypochlorite (NaOCl) and sodium chloride $$\textbf{Cl}_2 + 2\textbf{NaOH} \rightarrow \textbf{NaCl} + \textbf{NaOCl} + \textbf{H}_2\textbf{O}$$
2. With Hot Concentrated NaOH
   • Reaction: Produces sodium chlorate (NaClO₃) and sodium chloride (NaCl). $$3\textbf{Cl}_2 + 6\textbf{NaOH} \rightarrow 5\textbf{NaCl} + \textbf{NaClO}_3 + 3\textbf{H}_2\textbf{O}$$
3. With Excess NH₃
   • Reaction: Forms nitrogen and ammonium chloride (NH₄Cl). $$\textbf{Cl}_2 + 8\textbf{NH}_3 \rightarrow 6\textbf{NH}_4\textbf{Cl} + \textbf{N}_2$$
4. With Slaked Lime
   • Reaction: Produces calcium hypochlorite [Ca(OCl)₂], a component of bleaching powder. $$\textbf{Cl}_2 + \textbf{Ca}(\textbf{OH})_2 \rightarrow \textbf{Ca}(\textbf{OCl})_2 + \textbf{H}_2\textbf{O}$$
5. With Sodium Thiosulfate (Hypo)
   • Reaction: Chlorine oxidizes sodium thiosulfate to sodium tetrathionate (Na₂S₄O₆) and sodium chloride (NaCl). $$\textbf{Cl}_2 + \textbf{Na}_2\textbf{S}_2\textbf{O}_3 + \textbf{H}_2\textbf{O} \rightarrow \textbf{Na}_2\textbf{S}_4\textbf{O}_6 + 2\textbf{NaCl} + 2\textbf{H}_2\textbf{O}$$

Summary

The preparation and reactions of chlorine, from laboratory synthesis to industrial scale production via Deacon’s process, and its interactions with various substances, illustrate its versatility in chemical reactions and applications. Chlorine’s reactions with NaOH, NH₃, slaked lime, and sodium thiosulfate highlight its role in producing a range of compounds, from disinfectants to bleaching agents.

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LAQ-8 : How is chlorine prepared by electrolytic method? Explain its reaction with i) Iron ii) acidified FeSO₄ iii) Iodine iv) H₂S v) KI

Introduction

Chlorine, a highly reactive halogen, is extensively used in water treatment, disinfectants, and various chemical syntheses. It can be prepared through the electrolysis of brine (sodium chloride solution).

Preparation of Chlorine by Electrolytic Method

Electrolysis of Brine: Chlorine is produced at the anode during the electrolysis of aqueous sodium chloride (brine). The cathode reaction produces hydrogen gas, and the remaining solution contains sodium hydroxide. $$2\textbf{NaCl}(aq) + 2\textbf{H}_2\textbf{O}(l) \rightarrow 2\textbf{NaOH}(aq) + \textbf{H}_2(g) + \textbf{Cl}_2(g)$$

Reactions of Chlorine

1. With Iron (Fe)
   • Reaction: Chlorine reacts with iron to form iron(III) chloride (FeCl₃). $$2\textbf{Fe}(s) + 3\textbf{Cl}_2(g) \rightarrow 2\textbf{FeCl}_3(s)$$
2. With Acidified FeSO₄
   • Reaction: Chlorine oxidizes ferrous sulfate (FeSO₄) to ferric sulfate (Fe₂(SO₄)₃) and produces hydrogen chloride (HCl). $$2\textbf{FeSO}_4(aq) + \textbf{H}_2\textbf{SO}_4(aq) + \textbf{Cl}_2(g) \rightarrow \textbf{Fe}_2(\textbf{SO}_4)_3(aq) + 2\textbf{HCl}(aq)$$
3. With Iodine (I₂)
   • Reaction: Chlorine displaces iodine from iodine solution, demonstrating its higher reactivity. $$2\textbf{I}^-(aq) + \textbf{Cl}_2(g) \rightarrow \textbf{I}_2(s) + 2\textbf{Cl}^-(aq)$$
4. With Hydrogen Sulfide (H₂S)
   • Reaction: Chlorine reacts with hydrogen sulfide, converting it to sulfur and producing hydrochloric acid (HCl). $$\textbf{H}_2\textbf{S}(g) + \textbf{Cl}_2(g) \rightarrow 2\textbf{HCl}(aq) + \textbf{S}(s)$$
5. With Potassium Iodide (KI)
   • Reaction: Chlorine oxidizes potassium iodide to iodine, forming potassium chloride (KCl) and iodine in the process. $$2\textbf{KI}(aq) + \textbf{Cl}_2(g) \rightarrow 2\textbf{KCl}(aq) + \textbf{I}_2(s)$$

Summary

The electrolytic preparation of chlorine involves the decomposition of brine, yielding chlorine gas, hydrogen gas, and sodium hydroxide solution. Chlorine’s reactivity allows it to engage in various chemical reactions, such as forming compounds with metals like iron, oxidizing states in transition metals, displacing less reactive halogens like iodine, and reacting with hydrogen sulfide and potassium iodide. These reactions underscore chlorine’s role as a potent oxidizing agent in chemical processes.