Non-metal and their Compounds | Jamb Chemistry

Jamb chemistry key points on hydrogen; halogens; Hydrogen chloride and Hydrochloric acid etc

Hydrogen

Commercial Production of Hydrogen

1. Hydrogen is commercially produced using water gas ($CO + H_2$).
2. The water-gas shift reaction is used:
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   $CO + H_2O \rightarrow CO_2 + H_2$
3. Carbon dioxide is removed using solvents like potassium carbonate.
4. Steam reforming of methane is the most common method:
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   $CH_4 + H_2O \rightarrow CO + 3H_2$
5. Hydrogen is produced by electrolysis of water:
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   $2H_2O \rightarrow 2H_2 + O_2$
6. Cracking of petroleum fractions produces hydrogen by breaking long-chain hydrocarbons:
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   $C_nH_{2n+2} \rightarrow C_nH_{2n} + H_2$

Laboratory Preparation of Hydrogen

7. Hydrogen is prepared by reacting dilute acids with reactive metals:
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   $Zn + 2HCl \rightarrow ZnCl_2 + H_2$
8. Another method involves the reaction of water with alkali metals:
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   $2Na + 2H_2O \rightarrow 2NaOH + H_2$
9. Granulated zinc and dilute $H_2SO_4$ are commonly used in labs.
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Properties of Hydrogen

10. Hydrogen is a colorless, odorless, and tasteless gas.
11. It is the lightest element and less dense than air.
12. Hydrogen is highly flammable and burns with a pale blue flame.
13. It forms water upon combustion:
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    $2H_2 + O_2 \rightarrow 2H_2O$
14. Hydrogen is a reducing agent, reducing copper oxide to copper:
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    $CuO + H_2 \rightarrow Cu + H_2O$
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Uses of Hydrogen

15. Used in ammonia production in the Haber process:
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    $N_2 + 3H_2 \rightarrow 2NH_3$
16. Hydrogen is used in hydrogenation of oils to make margarine.
17. It serves as rocket fuel when combined with liquid oxygen.
18. Hydrogen is used in fuel cells for clean energy.
19. It acts as a reducing agent in metallurgy.
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Test for Hydrogen

20. Hydrogen burns with a “pop” sound when ignited in air, confirming its presence.
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Halogens

Overview of Halogens

21. Halogens are Group 17 elements: fluorine, chlorine, bromine, iodine, and astatine.
22. They are highly reactive nonmetals.
23. Halogens exist as diatomic molecules ($X_2$).
24. Reactivity decreases down the group: $F_2 > Cl_2 > Br_2 > I_2$.
25. Halogens are good oxidizing agents.
26. Halogens form salts (halides) with metals.
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Chlorine as a Representative Halogen

27. Chlorine is a greenish-yellow gas with a pungent smell.
28. It is denser than air and soluble in water.
29. Chlorine reacts with hydrogen to form hydrogen chloride:
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    $H_2 + Cl_2 \rightarrow 2HCl$
30. It bleaches organic dyes by oxidizing them.
31. Chlorine reacts with cold water to form a mixture of hydrochloric acid and hypochlorous acid:
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    $Cl_2 + H_2O \rightarrow HCl + HOCl$
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Laboratory Preparation of Chlorine

32. Chlorine is prepared by reacting $MnO_2$ with concentrated $HCl$:
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    $MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$
33. Alternatively, chlorine can be obtained by the electrolysis of brine.
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Industrial Preparation of Chlorine

34. Chlorine is industrially prepared by the electrolysis of brine ($NaCl$ solution).
35. At the anode, chloride ions are oxidized to chlorine gas:
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    $2Cl^- \rightarrow Cl_2 + 2e^-$
36. The chlor-alkali process also produces chlorine, hydrogen, and sodium hydroxide.
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Properties and Uses of Chlorine

Properties of Chlorine

37. Chlorine is a strong oxidizing agent.
38. It reacts with cold alkali to form hypochlorites:
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    $Cl_2 + 2NaOH \rightarrow NaCl + NaOCl + H_2O$
39. With hot alkali, it forms chlorates: $3Cl_2 + 6NaOH \rightarrow 5NaCl + NaClO_3 + 3H_2O$
40. Chlorine reacts with hydrocarbons to form chlorinated compounds (e.g., PVC).
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Uses of Chlorine

41. Chlorine is used for water disinfection.
42. It is essential in the production of bleaching powder.
43. Used in the manufacture of chlorinated organic compounds (e.g., CFCs).
44. Chlorine is used in the production of hydrochloric acid.
45. It is used as a bleaching agent in the textile and paper industries.
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Hydrogen Chloride and Hydrochloric Acid

Preparation of Hydrogen Chloride

46. Hydrogen chloride is prepared by heating sodium chloride with concentrated $H_2SO_4$:
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    $NaCl + H_2SO_4 \rightarrow NaHSO_4 + HCl$
47. $HCl$ gas is highly soluble in water, forming hydrochloric acid.
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Properties of Hydrogen Chloride

48. $HCl$ is a colorless gas with a pungent odor.
49. It fumes in moist air due to the formation of hydrochloric acid.
50. Hydrogen chloride is highly soluble in water, forming an acidic solution.
51. It reacts with ammonia to produce white ammonium chloride fumes:
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    $NH_3 + HCl \rightarrow NH_4Cl$
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Properties of Hydrochloric Acid

52. Hydrochloric acid is a strong acid.
53. It reacts with bases to form salts and water:
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    $HCl + NaOH \rightarrow NaCl + H_2O$
54. It reacts with carbonates and bicarbonates, releasing $CO_2$:
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    $HCl + NaHCO_3 \rightarrow NaCl + H_2O + CO_2$
55. $HCl$ reacts with metals like zinc to produce hydrogen gas:
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    $Zn + 2HCl \rightarrow ZnCl_2 + H_2$
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Uses of Hydrochloric Acid

56. Used in the pickling of steel to remove rust.
57. Essential in the production of chlorides and dyes.
58. It is used in laboratories for neutralization reactions.
59. Hydrochloric acid is a component of gastric juice, aiding digestion.
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Chlorides and Test for Chlorides

Overview of Chlorides

60. Chlorides are salts formed from the reaction of hydrochloric acid with bases or metals.
61. Common chlorides include $NaCl$, $KCl$, and $CaCl_2$.
62. Chlorides are ionic and typically soluble in water.
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Test for Chlorides

63. Chlorides react with silver nitrate solution in the presence of dilute nitric acid to form a white precipitate:
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    $Cl^- + Ag^+ \rightarrow AgCl {white precipitate}$
64. The precipitate dissolves in ammonia solution, confirming the presence of chloride ions.
65. The white precipitate darkens upon exposure to light.
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Significance of Halogens and Hydrogen Compounds

Applications in Daily Life

66. Chlorine disinfects drinking water, preventing waterborne diseases.
67. Hydrogen is a clean energy source for fuel cells.
68. Chlorides like $NaCl$ are essential for biological functions.
69. Bleaching powder is used in sanitation and public health.
70. Hydrochloric acid is crucial in the chemical industry.
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Environmental and Industrial Importance

71. Chlorine is used to treat wastewater.
72. Hydrogen's role in ammonia production is vital for fertilizers.
73. Halogens are used in flame retardants and pesticides.
74. Hydrogen chloride is used in the production of vinyl chloride (PVC).
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Experimental Procedures and Safety

Laboratory and Industrial Practices

75. Hydrogen preparation requires proper ventilation to prevent explosions.
76. Chlorine gas must be handled with care due to its toxicity.
77. $HCl$ solutions are corrosive and should be used with safety equipment.
78. Industrial production processes like electrolysis require precise control to optimize yields.
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Chemical Reactions

Specific Halogen Reactions

79. Chlorine reacts with metals to form chlorides:
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    $2Al + 3Cl_2 \rightarrow 2AlCl_3$
80. Hydrogen burns in oxygen, forming water:
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    $2H_2 + O_2 \rightarrow 2H_2O$
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Impact on Chemistry

81. Understanding reversible reactions is critical in halogen chemistry.
82. Reaction kinetics of halogens depend on their oxidizing strength.
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Miscellaneous

83. Hydrogen is the most abundant element in the universe.
84. Chlorine is the second most abundant halogen on Earth.
85. The solubility of halogens decreases down the group.
86. Hydrogen fuel cells produce only water as a byproduct, making them eco-friendly.
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Advanced Applications

87. Hydrogen isotopes like deuterium are used in nuclear reactors.
88. Chlorine compounds are critical in organic synthesis.
89. Hydrogen is used in space exploration as rocket fuel.
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Environmental Considerations

90. Chlorine production contributes to ozone depletion if not managed properly.
91. Hydrogen has potential for reducing carbon emissions.
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Significance in Chemistry

92. Chlorides are essential for studying ionic reactions.
93. Hydrogen is foundational in acid-base chemistry.
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Testing Knowledge

94. Identify the gas evolved when $Zn + HCl$ react.
95. What precipitate forms when $AgNO_3$ reacts with $NaCl$?
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Industrial Relevance

96. Chlorine enables mass production of household disinfectants.
97. Hydrogen is vital for the petroleum and chemical industries.
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Summarizing Insights

98. Halogens and hydrogen compounds are versatile in applications.
99. Their chemistry combines practicality with theoretical importance.
100. Mastery of these topics is crucial for advancements in both research and industry.
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Jamb chemistry Key points on Oxygen and Sulphur

Oxygen

Laboratory Preparation of Oxygen

1. Oxygen is prepared by heating potassium trioxochlorate(V) ($KClO_3$) with manganese(IV) oxide ($MnO_2$) as a catalyst:
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   $2KClO_3 \xrightarrow{heat, MnO}_2 2KCl + 3O_2$
2. The decomposition of hydrogen peroxide using $MnO_2$ also produces oxygen:
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   $2H_2O_2 \xrightarrow {MnO}_2 2H_2O + O_2$
3. Heating mercury(II) oxide releases oxygen:
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   $2HgO \xrightarrow{heat} 2Hg + O_2$
4. The gas is collected over water as it is slightly soluble in water.
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Properties of Oxygen

5. Oxygen is a colorless, odorless gas.
6. It supports combustion but is not combustible itself.
7. Oxygen is denser than air and slightly soluble in water.
8. It reacts with most elements to form oxides.
9. Oxygen reacts with hydrogen to form water:
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   $2H_2 + O_2 \rightarrow 2H_2O$
10. Combustion of carbon in oxygen produces carbon dioxide:
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    $C + O_2 \rightarrow CO_2$
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Uses of Oxygen

11. Used in oxy-acetylene torches for welding and cutting metals.
12. Essential for respiration in humans and animals.
13. Employed in medical oxygen therapy.
14. Used in the steel industry to remove impurities.
15. Oxygen is used in rocket propulsion systems.
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Commercial Production of Oxygen

16. Oxygen is commercially produced by fractional distillation of liquid air.
17. Nitrogen is removed at $-196^\circ{C}$, leaving oxygen, which boils at $-183^\circ{C}$.
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Oxides

Types of Oxides

18. Acidic oxides: Non-metal oxides that form acids in water (e.g., $CO_2$, $SO_2$).
19. Basic oxides: Metal oxides that form bases in water (e.g., $Na_2O$, $CaO$).
20. Amphoteric oxides: Oxides that react with both acids and bases (e.g., $Al_2O_3$, $ZnO$).
21. Neutral oxides: Oxides that do not react with acids or bases (e.g., $CO$, $N_2O$).
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Reactions of Oxides

Ozone (Trioxygen)

Ozone as an Allotrope of Oxygen

26. Ozone ($O_3$) is a pale blue gas with a characteristic smell.
27. It decomposes to oxygen, releasing energy:
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    $2O_3 \rightarrow 3O_2$
28. Ozone is a strong oxidizing agent.
29. It is formed by the action of electric discharge on oxygen:
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    $3O_2 \xrightarrow{electric discharge} 2O_3$
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Importance of Ozone

30. The ozone layer absorbs harmful ultraviolet radiation.
31. It protects living organisms from excessive UV exposure.
32. Depletion of the ozone layer leads to increased UV penetration.
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Sulphur

Allotropes of Sulphur

33. Sulphur exists as rhombic and monoclinic allotropes.
34. Rhombic sulphur is the most stable form at room temperature.
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Uses of Sulphur

35. Sulphur is used in the vulcanization of rubber.
36. It is a key ingredient in the manufacture of gunpowder and fireworks.
37. Sulphur is essential for producing sulfuric acid.
38. It acts as a fungicide and pesticide in agriculture.
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Sulphur(IV) Oxide

Preparation of

39. $SO_2$ is prepared by burning sulphur in oxygen:
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    $S + O_2 \rightarrow SO_2$
40. Heating metal sulfites with acids also produces $SO_2$:
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    $Na_2SO_3 + 2HCl \rightarrow 2NaCl + H_2O + SO_2$
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Properties of SO2

41. $SO_2$ is a colorless gas with a pungent smell.
42. It dissolves in water to form sulfurous acid:
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    $SO_2 + H_2O \rightarrow H_2SO_3$
43. It reduces $KMnO_4$ from purple to colorless:
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    $2KMnO_4 + 5SO_2 + 2H_2O \rightarrow 2MnSO_4 + K_2SO_4 + 2H_2SO_4$
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Uses of SO2

44. $SO_2$ is used as a bleaching agent in the paper industry.
45. It acts as a preservative for dried fruits.
46. It is employed in producing sulfuric acid.
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Reactions of SO2 with Alkalis

47. $SO_2$ reacts with sodium hydroxide to form sodium sulfite:
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    $SO_2 + 2NaOH \rightarrow Na_2SO_3 + H_2O$
48. With excess alkali, it forms sodium bisulfite:
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    $SO_2 + NaOH \rightarrow NaHSO_3$
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Trioxosulphate(IV) Acid and Its Salts

49. Trioxosulphate(IV) acid ($H_2SO_3$) is weak and unstable.
50. It decomposes into $SO_2$ and $H_2O$.
51. Acids react with sulfites to release $SO_2$:
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    $Na_2SO_3 + 2HCl \rightarrow 2NaCl + SO_2 + H_2O$
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Tetraoxosulphate(VI) Acid

Commercial Preparation

52. $H_2SO_4$ is prepared by the contact process.
53. Sulphur is burned to form $SO_2$:
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    $S + O_2 \rightarrow SO_2$
54. $SO_2$ is oxidized to $SO_3$ using $V_2O_5$ as a catalyst: $2SO_2 + O_2 \leftrightarrow 2SO_3$
55. $SO_3$ dissolves in $H_2SO_4$ to form oleum:
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    $SO_3 + H_2SO_4 \rightarrow H_2S_2O_7$
56. Oleum is diluted to produce $H_2SO_4$:
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    $H_2S_2O_7 + H_2O \rightarrow 2H_2SO_4$
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Properties of H2SO4

57. $H_2SO_4$ is a strong acid.
58. It is a dehydrating agent, removing water from organic compounds.
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Uses of H2So4

59. Used in fertilizer production (e.g., superphosphates).
60. Employed in the manufacture of detergents.
61. It is used as an electrolyte in car batteries.
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Test for SO4(2-)

62. Sulfate ions react with barium chloride in the presence of dilute $HCl$ to form a white precipitate:
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    $SO_4^{2-} + Ba^{2+} \rightarrow BaSO_4$
63. The precipitate is insoluble in acids, confirming the presence of sulfate ions.
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Applications and Significance

64. Oxygen is essential for life and combustion processes.
65. Ozone protects living organisms by filtering UV radiation.
66. Sulphur compounds are vital for fertilizers and industrial chemicals.
67. $SO_2$ is significant in producing sulfuric acid.
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Environmental Impact

68. Excessive $SO_2$ emissions lead to acid rain.
69. Ozone layer depletion harms ecosystems and human health.
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Laboratory Procedures

70. Careful handling of $SO_2$ and $H_2SO_4$ is crucial due to toxicity and corrosiveness.
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Summary of Reactions

71. Oxygen rekindles a glowing splint.
72. Sulfate ions form a white precipitate with $BaCl_2$.
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Miscellaneous Facts

73. Oxygen is the second most abundant gas in the Earth's atmosphere.
74. Sulphur is widely found in nature as sulfide and sulfate minerals.
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Industrial Relevance

75. Oxygen is critical in steel production.
76. $H_2SO_4$ is a benchmark for industrial capacity.
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Advanced Applications

77. Sulphur isotopes are used in tracer studies.
78. Ozone disinfects air and water in small quantities.
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Conclusion

79. Oxygen, ozone, and sulphur compounds have diverse industrial and environmental roles.
80. Mastery of these topics is essential for both academic and practical applications.
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Jamb chemistry Key points on Hydrogen sulphide

Preparation and Properties of Hydrogen Sulphide

1. Hydrogen sulphide is prepared by reacting iron(II) sulfide ($FeS$) with dilute hydrochloric acid:
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   $FeS + 2HCl \rightarrow FeCl_2 + H_2S$
2. It is a colorless gas with a characteristic rotten egg smell.
3. $H_2S$ is slightly soluble in water, forming a weak acidic solution:
   
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   $H_2S \leftrightarrow H^+ + HS^-$
4. As a weak acid, it reacts with strong bases to form sulfides: $H_2S + 2NaOH \rightarrow Na_2S + 2H_2O$
5. $H_2S$ is a reducing agent, reducing iodine to hydrogen iodide:
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   $H_2S + I_2 \rightarrow 2HI + S$
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Reducing and Precipitating Agents

6. $H_2S$ reduces sulfur dioxide to elemental sulfur:
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   $2H_2S + SO_2 \rightarrow 3S + 2H_2O$
7. It precipitates metal ions as insoluble sulfides (e.g., $Pb^{2+}$ forms $PbS$).
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Test for Sulphide ions

8. Sulfide ions react with dilute acids, releasing $H_2S$ gas, which smells like rotten eggs.
9. Lead acetate reacts with $S^{2-}$ to form a black precipitate of $PbS$:
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   $S^{2-} + Pb^{2+} \rightarrow PbS$
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Laboratory Preparation of Nitrogen

10. Nitrogen is prepared by heating ammonium nitrite:
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    $NH_4NO_2 \rightarrow N_2 + 2H_2O$
11. The gas is collected over water and purified using sodium hydroxide.
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Production of Nitrogen from Liquid Air

12. Liquid air is subjected to fractional distillation.
13. Nitrogen, with a boiling point of $-196^\circ{C}$, boils off first, leaving oxygen behind.
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Ammonia

Laboratory Preparation of Ammonia

14. Ammonia is prepared by heating ammonium chloride with sodium hydroxide:
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    $NH_4Cl + NaOH \rightarrow NH_3 + NaCl + H_2O$
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Industrial Preparation (Haber Process)

15. The Haber process synthesizes ammonia from nitrogen and hydrogen:
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    $N_2 + 3H_2 \leftrightarrow 2NH_3$
16. The reaction occurs at 450°C, 200 atm pressure, and uses an iron catalyst.
17. High pressure favors ammonia formation due to fewer gas molecules in the products.
18. Moderate temperatures optimize the yield and rate of reaction.
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Properties of Ammonia

19. Ammonia is a colorless gas with a pungent smell.
20. It is highly soluble in water, forming ammonium hydroxide:
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    $NH_3 + H_2O \leftrightarrow NH_4^+ + OH^-$
21. Ammonia neutralizes acids to form salts:
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    $NH_3 + HCl \rightarrow NH_4Cl$
22. It reduces copper(II) oxide to copper:
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    $ 3CuO + 2NH_3 \rightarrow 3Cu + N_2 + 3H_2O
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Uses of Ammonia

23. Ammonia is used in fertilizers (e.g., ammonium nitrate and urea).
24. It is a refrigerant in cooling systems.
25. Ammonia is essential in the production of nitric acid.
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Ammonium Salts and Their Uses

26. Ammonium sulfate ($(NH_4)_2SO_4$) is a fertilizer.
27. Ammonium chloride ($NH_4Cl$) is used in dry cell batteries.
28. Ammonium nitrate ($NH_4NO_3$) is used in explosives and fertilizers.
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Oxidation of Ammonia

29. Ammonia is oxidized to nitric oxide in the Ostwald process:
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    $4NH_3 + 5O_2 \rightarrow 4NO + 6H_2O$
30. Nitric oxide is further oxidized to nitrogen dioxide:
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    $2NO + O_2 \rightarrow 2NO_2$
31. Nitrogen dioxide reacts with water and oxygen to form nitric acid:
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    $4NO_2 + 2H_2O + O_2 \rightarrow 4HNO_3$
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Test for Ammonium ions

32. Adding sodium hydroxide to ammonium salts releases ammonia gas:
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    $NH_4Cl + NaOH \rightarrow NH_3 + NaCl + H_2O$
33. Ammonia gas turns moist red litmus paper blue.
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Trioxonitrate(V) Acid

Laboratory Preparation

34. Nitric acid is prepared by heating sodium nitrate with concentrated sulfuric acid:
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    $NaNO_3 + H_2SO_4 \rightarrow NaHSO_4 + HNO_3$
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Trioxonitrate(V) Salts

Action of Heat on Trioxonitrate(V) Salts

35. Sodium nitrate decomposes to form sodium nitrite and oxygen:
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    $2NaNO_3 \xrightarrow{heat} 2NaNO_2 + O_2$
36. Ammonium nitrate decomposes explosively to form nitrous oxide and water:
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    $NH_4NO_3 \xrightarrow{heat} N_2O + 2H_2O$
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Uses of Trioxonitrate(V) Salts

37. Potassium nitrate ($KNO_3$) is used in fertilizers and gunpowder.
38. Sodium nitrate ($NaNO_3$) is used as a preservative in food.
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Test for Nitrate ions

39. Nitrate ions produce brown fumes of nitrogen dioxide when heated with sulfuric acid and copper turnings:
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    $2NO_3^- + 4H^+ + 2Cu \rightarrow 2NO_2 + 2H_2O + 2Cu^{2+}$
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Oxides of Nitrogen

40. Nitrogen forms dinitrogen oxide ($N_2O$), nitric oxide ($NO$), and nitrogen dioxide ($NO_2$).
41. $N_2O$ is a colorless gas with a sweet smell, used as an anesthetic.
42. $NO$ is a colorless gas that turns brown in air due to oxidation to $NO_2$.
43. $NO_2$ is a reddish-brown toxic gas.
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The Nitrogen Cycle

44. The nitrogen cycle describes the movement of nitrogen through the environment.
45. Atmospheric nitrogen is fixed by bacteria into ammonia.
46. Nitrifying bacteria convert ammonia to nitrites:
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    $NH_3 \rightarrow NO_2^-$
47. Nitrites are further oxidized to nitrates:
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    $NO_2^- \rightarrow NO_3^-$
48. Plants absorb nitrates from the soil for protein synthesis.
49. Denitrifying bacteria convert nitrates back to nitrogen gas:
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    $NO_3^- \rightarrow N_2$
50. Decomposition of organic matter returns nitrogen to the soil as ammonia.
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Jamb chemistry Key points on Carbon

Uses and Properties of Allotropes

Definition and Overview

1. Allotropes are different forms of the same element in the same physical state with distinct properties.
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Allotropes of Carbon

2. Carbon exists as diamond, graphite, fullerenes, and amorphous carbon.
3. Diamond: Hardest natural substance, excellent insulator, high refractive index.
4. Graphite: Soft, slippery, good conductor of electricity, high melting point.
5. Fullerenes: Hollow spherical or tubular molecules with unique structural properties.
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Uses of Carbon Allotropes

6. Diamond: Used in cutting tools, jewelry, and industrial abrasives.
7. Graphite: Used in lubricants, pencils, electrodes, and nuclear reactors.
8. Fullerenes: Used in drug delivery, nanotechnology, and superconductors.
9. Amorphous carbon (e.g., charcoal, lampblack): Used in filtration, pigments, and fuel.
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Carbon(IV) Oxide

Laboratory Preparation

10. Carbon(IV) oxide is prepared by reacting dilute acids with trioxocarbonate salts: $CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2$
11. The gas is collected by downward delivery due to its high density.
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Properties of Carbon(IV) Oxide

12. It is a colorless, odorless gas.
13. It is denser than air and soluble in water, forming carbonic acid:
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    $CO_2 + H_2O \leftrightarrow H_2CO_3$
14. It extinguishes flames and does not support combustion.
15. $CO_2$ turns limewater milky due to the formation of calcium carbonate:
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    $CO_2 + Ca(OH)_2 \rightarrow CaCO_3 + H_2O$
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Uses of Carbon(IV) Oxide

16. Used in fire extinguishers.
17. Essential in photosynthesis by plants.
18. Used to carbonate soft drinks.
19. It is a refrigerant in cooling systems (dry ice).
20. Used in enhanced oil recovery for extraction processes.
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Trioxocarbonate(IV) salts

Action of Heat on Trioxocarbonate(IV) Salts

21. Heating metal trioxocarbonate(IV) salts releases carbon(IV) oxide:
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    $CaCO_3 \xrightarrow{\text{heat}} CaO + CO_2$
22. Group I carbonates (except lithium carbonate) do not decompose when heated.
23. Metal carbonates decompose into oxides and carbon dioxide.
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Test for carbonate ions

24. Adding dilute acids to carbonates produces carbon(IV) oxide:
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    $CO_3^{2-} + 2H^+ \rightarrow CO_2 + H_2O$
25. Carbon(IV) oxide turns limewater milky, confirming $CO_3^{2-}$.
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Carbon(II) Oxide (CO)

Laboratory Preparation

26. Carbon(II) oxide is prepared by heating oxalic acid with concentrated sulfuric acid:
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    H_2C_2O_4 \xrightarrow{\text{H_2SO_4}} CO + CO_2 + H_2O
27. Carbon(IV) oxide is removed by passing the gas through sodium hydroxide solution.
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Properties of Carbon(II) Oxide

28. Carbon(II) oxide is a colorless, odorless, and highly toxic gas.
29. It is slightly less dense than air.
30. It is a reducing agent, reducing $Fe_2O_3$ to iron:
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    $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$
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Uses of Carbon(II) Oxide

31. Used in the production of methanol.
32. Essential in the synthesis of phosgene and organic compounds.
33. Used as a reducing agent in metallurgy.
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Effects of Carbon(II) Oxide in Blood

34. Carbon(II) oxide binds to hemoglobin to form carboxyhemoglobin.
35. This prevents oxygen from binding to hemoglobin, reducing oxygen transport.
36. High exposure causes suffocation, dizziness, and even death.
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Sources of Carbon(II) Oxide

37. Charcoal fires produce $CO$ when air supply is limited.
38. Vehicle exhaust fumes are a significant source of $CO$.
39. Incomplete combustion of hydrocarbons releases $CO$.
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Destructive Distillation

Products from Destructive Distillation of Wood

40. Destructive distillation of wood produces charcoal, wood tar, and wood vinegar.
41. Wood gas (a mixture of $H_2, CH_4, CO,$ and $CO_2$) is also released.
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Products from Destructive Distillation of Coal

42. Destructive distillation of coal produces coke, coal tar, ammoniacal liquor, and coal gas.
43. Coke is a solid residue used as fuel and in steelmaking.
44. Coal tar is a source of aromatic hydrocarbons for chemicals and dyes.
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Gasification and Uses of Coke

Gasification of Coke

45. Coke reacts with steam to produce synthesis gas ($CO + H_2$):
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    $C + H_2O \rightarrow CO + H_2$
46. The process is carried out in gasifiers at high temperatures.
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    ####### Uses of Coke
47. Used as a reducing agent in blast furnaces for iron extraction.
48. It is a fuel in industrial and domestic settings.
49. Coke is used in the production of water gas.
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Synthesis Gas

Manufacture of Synthesis Gas

50. Synthesis gas is produced by the steam reforming of methane:
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    $CH_4 + H_2O \rightarrow CO + 3H_2$
51. It can also be produced via partial oxidation of hydrocarbons:
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    $2CH_4 + O_2 \rightarrow 2CO + 4H_2$

Uses of Synthesis Gas

52. Used in the production of methanol:
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    $CO + 2H_2 \rightarrow CH_3OH$
53. Essential in ammonia synthesis via the Haber process.
54. Used to produce synthetic hydrocarbons via the Fischer-Tropsch process.
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Summary

55. Carbon allotropes have diverse physical properties and industrial applications.
56. Carbon(IV) oxide is non-toxic and widely used, while carbon(II) oxide is toxic and primarily a reducing agent.
57. The decomposition of trioxocarbonate salts produces ( CO_2 ) and oxides.
58. Destructive distillation of wood and coal provides valuable chemicals and fuels.
59. Synthesis gas is a cornerstone of modern chemical manufacturing.

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