Metal and their Compounds | Jamb Chemistry

Jamb chemistry key points on General properties of metals; Alkali metals e.g. sodium

General Properties of Metals

1. Metals are good conductors of heat and electricity due to free electrons.
2. Most metals have high melting and boiling points.
3. Metals are malleable; they can be hammered into thin sheets.
4. They are ductile, meaning they can be drawn into wires.
5. Metals are lustrous (shiny) when freshly cut or polished.
6. They are sonorous, producing a ringing sound when struck.
7. Metals are typically solid at room temperature, except mercury.
8. Metals have high densities, though alkali metals have lower densities.
9. Metals form cations by losing electrons, making them electropositive.
10. They react with oxygen to form basic oxides.
11. Metals displace hydrogen from dilute acids to form salts and hydrogen gas.
12. Most metals form alloys with other metals or non-metals.
13. Metals are strong and have high tensile strength.
14. They form ionic bonds with non-metals due to their low electronegativity.
    
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Alkali Metals: Sodium

15. Alkali metals are highly reactive and soft, with low melting points.
16. Sodium is a silvery-white, soft metal that can be cut with a knife.
17. It is stored under oil because it reacts with moisture and oxygen.
18. Sodium reacts vigorously with water, forming sodium hydroxide and hydrogen:
    
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    $2Na + 2H_2O \rightarrow 2NaOH + H_2$
19. Sodium burns with a yellow flame in air, forming sodium oxide:
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    $4Na + O_2 \rightarrow 2Na_2O$
20. Sodium reacts with chlorine to form sodium chloride ($NaCl$).
21. Sodium compounds, such as sodium hydroxide, sodium carbonate, and sodium chloride, are industrially important.
    
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Production of Sodium Hydroxide by Electrolysis of Brine

22. Sodium hydroxide ($NaOH$) is produced by the electrolysis of brine ($NaCl$ solution).
23. In the process, brine is electrolyzed in a diaphragm cell.
24. At the cathode, hydrogen gas is produced:
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    $2H_2O + 2e^- \rightarrow H_2 + 2OH^-$
25. At the anode, chlorine gas is liberated:
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    $2Cl^- \rightarrow Cl_2 + 2e^-$
26. Sodium ions remain in solution and combine with hydroxide ions to form sodium hydroxide:
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    $Na^+ + OH^- \rightarrow NaOH$
27. The by-products are hydrogen and chlorine gases.
28. The resulting solution is concentrated to obtain solid $NaOH$.
29. Sodium hydroxide is widely used in soap making and paper production.
    
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Action of Sodium Hydroxide on Aluminium, Zinc, and Lead Ions

30. Sodium hydroxide reacts with aluminium ions to form a white precipitate of aluminium hydroxide:
    
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    $Al^{3+} + 3OH^- \rightarrow Al(OH)_3$
31. The precipitate dissolves in excess $NaOH$, forming a soluble aluminate complex:
    
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    $Al(OH)_3 + NaOH \rightarrow Na[Al(OH)_4]$
32. Zinc ions react with $NaOH$ to form a white precipitate of zinc hydroxide:
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    $Zn^{2+} + 2OH^- \rightarrow Zn(OH)_2$
33. Zinc hydroxide dissolves in excess $NaOH$:
    
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    $Zn(OH)_2 + 2NaOH \rightarrow Na_2[Zn(OH)_4]$
34. Lead ions form a white precipitate of lead(II) hydroxide with sodium hydroxide:
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    $Pb^{2+} + 2OH^- \rightarrow Pb(OH)_2$
35. Lead(II) hydroxide does not dissolve in excess sodium hydroxide.
    
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Precipitation of Metallic Hydroxides

36. Sodium hydroxide is used to test for metal ions via precipitation reactions.
37. Iron(II) ions form a green precipitate of iron(II) hydroxide:
    
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    $Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2$
38. Iron(III) ions form a reddish-brown precipitate of iron(III) hydroxide:
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    $Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3$
39. Copper(II) ions form a blue precipitate of copper(II) hydroxide:
    
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    $Cu^{2+} + 2OH^- \rightarrow Cu(OH)_2$
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Sodium Trioxocarbonate(IV) and Sodium Hydrogen Trioxocarbonate(IV)

Production of Sodium Trioxocarbonate(IV)

40. Sodium trioxocarbonate(IV) ($Na_2CO_3$) is produced via the Solvay process.
41. In this process, ammonia and carbon dioxide react with sodium chloride:
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    $NaCl + NH_3 + CO_2 + H_2O \rightarrow NaHCO_3 + NH_4Cl$
42. Sodium bicarbonate ($NaHCO_3$) is heated to form sodium carbonate:
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    $2NaHCO_3 \xrightarrow{heat} Na_2CO_3 + CO_2 + H_2O$
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Properties of Sodium Carbonates

43. Sodium carbonate is a white crystalline solid soluble in water.
44. It forms an alkaline solution in water.
45. Sodium hydrogen trioxocarbonate(IV) ($NaHCO_3$) is a weak base and decomposes on heating:
    
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    $2NaHCO_3 \xrightarrow{heat} Na_2CO_3 + CO_2 + H_2O$
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Sodium Carbonate in the Manufacture of Glass

46. Sodium carbonate is a key raw material in glass production.
47. It reduces the melting point of silica ($SiO_2$).
48. It reacts with silica to form sodium silicate:
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    $Na_2CO_3 + SiO_2 \rightarrow Na_2SiO_3 + CO_2$
49. Calcium carbonate is added to improve glass strength and durability.
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    # Sodium Chloride

Occurrence in Sea Water

50. Sodium chloride ($NaCl$) occurs naturally in sea water, salt lakes, and rock salt.
51. It constitutes about 2.7% of sea water by mass.
    
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Uses of Sodium Chloride

52. It is used as table salt for seasoning food.
53. Sodium chloride is used in the electrolysis process to produce sodium hydroxide, chlorine, and hydrogen.
54. It is used in the production of soda ash ($Na_2CO_3$) via the Solvay process.
55. Sodium chloride is spread on roads to melt ice during winter.
56. It is used in the preservation of food (e.g., fish and meat).
    
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Importance of Sea Water

57. Sea water is a major source of common salt ($NaCl$).
58. It contains dissolved minerals such as magnesium, bromine, and iodine.
59. Sea water supports the fishing industry, providing food and employment.
60. It is a source of desalinated water for drinking and agriculture.
61. Sea water supports marine biodiversity and ecosystems.
62. It contributes to tourism and recreational activities.
    
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Recovery of Sodium Chloride

63. Sodium chloride is recovered from sea water through solar evaporation.
64. Sea water is collected in shallow evaporation ponds.
65. Sunlight evaporates the water, leaving behind crystalline salt.
66. Impurities are removed through washing and refining.
67. Recovered salt is used for industrial and domestic purposes.
    
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Summary

68. Metals have common physical and chemical properties such as malleability and conductivity.
69. Sodium is a highly reactive alkali metal with industrial significance.
70. Sodium hydroxide is produced via the electrolysis of brine.
71. Sodium hydroxide reacts differently with aluminium, zinc, and lead ions.
72. Sodium carbonate and bicarbonate are essential industrial chemicals.
73. Sodium carbonate is crucial in glass manufacturing.
74. Sodium chloride is abundant in sea water and has wide industrial and domestic uses.
75. Sea water provides economic benefits through salt production, fishing, and desalination.
    
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Jamb chemistry Key points on Alkaline-earth metals; Aluminium; Tin

Alkaline-Earth Metals: Calcium

1. Calcium is an alkaline-earth metal found in Group 2 of the periodic table.
2. It is a reactive, silvery-white metal that forms a protective oxide coating in air.
3. Calcium has a relatively low density (1.55 g/cm³).
4. It reacts with water to form calcium hydroxide and hydrogen gas:
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   $Ca + 2H_2O \rightarrow Ca(OH)_2 + H_2$
5. Calcium reacts with oxygen to form calcium oxide:
   
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   $2Ca + O_2 \rightarrow 2CaO$
6. Calcium is essential for biological processes, including bone and teeth formation.
7. It forms ionic compounds, e.g., calcium chloride ($CaCl_2$) and calcium trioxocarbonate(IV) ($CaCO_3$).
   
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Calcium Oxide, Calcium Hydroxide, and Calcium Trioxocarbonate(IV)

8. *Calcium oxide ($CaO )*$ is commonly known as quicklime.
9. It is produced by heating calcium carbonate:
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   $CaCO_3 \xrightarrow{heat} CaO + CO_2$
10. Calcium oxide reacts with water to form calcium hydroxide ($Ca(OH)_2$):
    
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    $CaO + H_2O \rightarrow Ca(OH)_2$
11. Calcium hydroxide ($Ca(OH)_2$), or slaked lime, is a white, alkaline powder.
12. It is sparingly soluble in water to form limewater.
13. Limewater turns milky when exposed to carbon dioxide:
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    $Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O$
14. Calcium trioxocarbonate(IV) ($CaCO_3$) occurs naturally as limestone, chalk, and marble.
15. It is insoluble in water but reacts with acids to produce carbon dioxide:
    
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    $CaCO_3 + 2HCl \rightarrow CaCl_2 + CO_2 + H_2O$
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Properties and Uses of Calcium Hydroxide and Calcium Trioxocarbonate(IV)

16. Calcium hydroxide is used in neutralizing acidic soils in agriculture.
17. It is used in water treatment to remove impurities.
18. Calcium hydroxide is used in the production of mortar and whitewash.
19. Calcium trioxocarbonate(IV) is a raw material for cement production.
20. It is used in the manufacture of lime, glass, and quicklime.
21. Calcium carbonate is an antacid for treating indigestion.
    
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Preparation of Calcium Oxide from Sea Shells

22. Sea shells are rich in calcium carbonate ($CaCO_3$).
23. They are heated in a kiln to decompose calcium carbonate into calcium oxide and carbon dioxide:
    
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    $CaCO_3 \xrightarrow{heat} CaO + CO_2$
24. The calcium oxide (quicklime) is collected and used in industries.
    
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Chemical Composition of Cement and the Setting of Mortar

25. Cement is made by heating limestone ($CaCO_3$) and clay to produce clinker.
26. Clinker contains compounds like tricalcium silicate ($3CaO·SiO_2$) and dicalcium silicate ($2CaO·SiO_2$).
27. Gypsum ($CaSO_4·2H_2O$) is added to control the setting time.
28. When cement is mixed with water, hydration occurs, forming a hard solid.
29. Mortar is a mixture of cement, sand, and water.
30. The setting of mortar involves the hardening of hydrated calcium silicates.
    
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Test for Calcium ions

31. Calcium ions form a white precipitate with sodium hydroxide:
    
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    $Ca^{2+} + 2OH^- \rightarrow Ca(OH)_2$
32. Calcium ions give a brick-red flame in a flame test.
    
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Aluminium

33. Aluminium is a lightweight, silvery-white metal found in Group 13 of the periodic table.
34. It has a low density (2.7 g/cm³) and is a good conductor of electricity and heat.
35. It resists corrosion due to the formation of a thin oxide layer.
36. Aluminium is amphoteric, reacting with both acids and bases:
    
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    $Al + 3HCl \rightarrow AlCl_3 + 3H_2$ $2Al + 2NaOH + 6H_2O \rightarrow 2Na[Al(OH)_4] + 3H_2$
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Purification of Bauxite

37. Bauxite ($Al_2O_3·2H_2O$) is the main ore of aluminium.
38. Impurities like iron oxides and silica are removed by the Bayer process.
39. Bauxite is digested with concentrated sodium hydroxide solution.
40. Aluminium oxide dissolves, leaving impurities as red mud.
41. The solution is filtered, and aluminium hydroxide is precipitated by cooling.
42. Aluminium hydroxide is calcined (heated) to produce pure aluminium oxide ($Al_2O_3$).
    
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Electrolytic Extraction of Aluminium

43. Aluminium is extracted by the electrolysis of molten aluminium oxide dissolved in cryolite.
44. The process occurs in a steel tank lined with carbon (cathode).
45. At the cathode, aluminium is deposited:
    
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    $Al^{3+} + 3e^- \rightarrow Al$
46. At the anode, oxygen gas is released, which reacts with carbon electrodes:
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    $2O^{2-} \rightarrow O_2 + 4e^-$
47. Cryolite lowers the melting point of aluminium oxide and improves conductivity.
    
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Properties and Uses of Aluminium

48. Aluminium is lightweight, strong, and corrosion-resistant.
49. It is used in making aircraft, cars, and packaging materials (e.g., foil).
50. It is used in power lines due to its electrical conductivity.
51. Aluminium alloys, like duralumin, are used in construction.
52. Aluminium oxide is used as an abrasive and in ceramics.
53. Aluminium hydroxide is used in antacids and water purification.
    
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Test for Aluminium ions

54. Aluminium ions form a white precipitate with sodium hydroxide:
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    $Al^{3+} + 3OH^- \rightarrow Al(OH)_3$
55. The precipitate dissolves in excess $NaOH$:
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    $Al(OH)_3 + NaOH \rightarrow Na[Al(OH)_4]$
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Tin

56. Tin is a soft, silvery-white metal found in Group 14 of the periodic table.
57. It has a low melting point (232°C) and resists corrosion.
58. Tin exists in two allotropes: white tin (stable) and grey tin (unstable at low temperatures).
    
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Extraction of Tin from Its Ores

59. The main ore of tin is cassiterite ($SnO_2$).
60. The ore is first concentrated by washing and gravity separation.
61. The concentrated ore is roasted to remove impurities like sulfur and arsenic.
62. Tin(IV) oxide is reduced with carbon in a furnace:
    
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    $SnO_2 + 2C \rightarrow Sn + 2CO$
63. Impurities are removed by liquation or electrolytic refining.
    
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Properties and Uses of Tin

64. Tin is malleable, ductile, and resistant to corrosion.
65. It is used to coat steel to prevent rusting (tin plating).
66. Tin alloys, like bronze (tin and copper), are used in statues and tools.
67. Tin is used in soldering as part of soft solder (tin and lead).
68. It is used in food cans due to its non-toxic nature.
    
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Summary

69. Calcium is an essential alkaline-earth metal with industrial uses in lime and cement production.
70. Calcium hydroxide and calcium carbonate have versatile applications.
71. Bauxite is the primary source of aluminium, extracted via electrolysis.
72. Aluminium is widely used due to its strength, lightness, and conductivity.
73. Tin is extracted from cassiterite and used in coating, alloys, and soldering.
    
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Jamb chemistry Key points on Metals of the first transition series; Iron

General Introduction to the First Transition Series

1. The first transition series consists of elements with atomic numbers 21 to 30: $Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu,$ and $Zn$.
2. These elements have partially filled $d$-orbitals, which give them unique properties.
3. They are located in Period 4 and the $d$-block of the periodic table.
   
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Characteristic Properties of Electron Configuration

4. Transition metals have electron configurations involving the filling of $3d$-orbitals.
5. General configuration: $[Ar] 3d^{1-10} 4s^2$.
6. Scandium ($Sc$) has the simplest configuration: $[Ar] 3d^1 4s^2$.
7. Zinc ($Zn$) has a completely filled $d$-orbital: $[Ar] 3d^{10} 4s^2$.
8. Electron configurations often exhibit exceptions due to extra stability of half-filled and fully filled $d$-subshells.
9. Example: Chromium ($Cr$) has $[Ar] 3d^5 4s^1$ instead of $[Ar] 3d^4 4s^2$.
10. Copper ($Cu$) has $[Ar] 3d^{10} 4s^1$ instead of $[Ar] 3d^9 4s^2$.
    
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Characteristic Properties of Oxidation States

11. Transition metals exhibit variable oxidation states due to the availability of $3d$ and $4s$ electrons for bonding.
12. The oxidation states increase across the series and then decrease.
13. Scandium ($Sc$) exhibits only +3 oxidation state.
14. Vanadium ($V$) exhibits oxidation states +2, +3, +4, and +5.
15. Manganese ($Mn$) shows oxidation states from +2 to +7.
16. Iron ($Fe$) commonly shows +2 and +3 oxidation states.
17. Copper ($Cu$) exhibits +1 and +2 oxidation states.
18. Zinc ($Zn$) shows only the +2 oxidation state due to a filled $3d$-orbital.
19. Higher oxidation states are often stabilized in oxides or fluorides.
20. Lower oxidation states tend to form ionic compounds, while higher ones form covalent compounds.
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Characteristic Properties of Complex Ion Formation

21. Transition metals readily form complex ions due to small size and high charge density.
22. Complex ions contain a central metal ion surrounded by ligands.
23. Ligands donate lone pairs of electrons to form coordinate covalent bonds.
24. Examples of ligands: $NH_3, H_2O, Cl^-, CN^-,$ and $OH^-$.
25. Iron forms complexes such as $[Fe(H_2O)_6]^{2+}$ and $[Fe(CN)_6]^{3-}$.
26. Copper forms the complex ion $[Cu(NH_3)_4]^{2+}$ with ammonia.
27. Coordination numbers (number of ligands) are commonly 4 or 6 for transition metals.
28. Complex ion formation stabilizes metal ions in solutions.
    
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Characteristic Properties of Formation of Coloured Ions

29. Transition metals form coloured compounds due to $d-d$ electron transitions.
30. Electrons absorb visible light to move between split $d$-orbitals.
31. The colour depends on the energy gap between the split $d$-orbitals.
32. $Fe^{2+}$ ions are pale green, while $Fe^{3+}$ ions are yellow-brown.
33. $Cu^{2+}$ ions are blue in aqueous solution.
34. $Cr^{3+}$ ions are green, and $Mn^{2+}$ ions are pink.
35. Colour intensity increases with increasing oxidation state.
36. Zinc and scandium ions are colourless due to fully filled or empty $d$-orbitals.
    
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Characteristic Properties of Catalysis

37. Transition metals and their compounds act as catalysts in many reactions.
38. They provide a surface for reactants to adsorb and react.
39. Transition metals can change oxidation states, providing alternative reaction pathways.
40. Iron acts as a catalyst in the Haber process for ammonia production:
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    $N_2 + 3H_2 \rightarrow 2NH_3$
41. Vanadium(V) oxide ($V_2O_5$) catalyzes the oxidation of sulfur dioxide to sulfur trioxide in the Contact process:
    
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    $2SO_2 + O_2 \rightarrow 2SO_3$
42. Nickel acts as a catalyst in the hydrogenation of vegetable oils.
43. Manganese(IV) oxide catalyzes the decomposition of hydrogen peroxide:
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    $2H_2O_2 \rightarrow 2H_2O + O_2$
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Iron

Extraction of Iron

44. Iron is extracted from its oxides, primarily hematite ($Fe_2O_3$) and magnetite ($Fe_3O_4$).
45. Iron sulfide ($FeS_2$) is roasted to produce iron oxide:
    
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    $4FeS_2 + 11O_2 \rightarrow 2Fe_2O_3 + 8SO_2$
46. Hematite is reduced in a blast furnace using carbon and carbon monoxide.
47. The main reactions in the blast furnace are:
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    $C + O_2 \rightarrow CO_2$
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    $CO_2 + C \rightarrow 2CO$
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    $Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2$
48. Limestone ($CaCO_3$) is added to remove impurities, forming slag:
    
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    $CaCO_3 \rightarrow CaO + CO_2$
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    $CaO + SiO_2 \rightarrow CaSiO_3$
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Different Forms of Iron and Their Properties

49. Pig iron: High carbon content (4-5%), brittle and hard.
50. Cast iron: Slightly less carbon than pig iron, used for engine blocks.
51. Wrought iron: Low carbon content (less than0.1%), malleable and tough.
52. Steel: An alloy of iron with controlled carbon content (less than 2%).
    
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Advantages of Steel Over Iron

53. Steel is stronger and more durable than pure iron.
54. It is less brittle compared to cast iron.
55. Steel can be alloyed with other elements to enhance properties (e.g., stainless steel).
56. Steel resists corrosion better than untreated iron.
    
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**Test for Ferrous ions and Ferric ions

57. Test for $Fe^{2+}$: Add sodium hydroxide; a green precipitate of $Fe(OH)_2$ forms:
    
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    $Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2$
58. The precipitate turns brown in air due to oxidation to $Fe(OH)_3$.
59. Test for $Fe^{3+}$: Add sodium hydroxide; a brown precipitate of $Fe(OH)_3$ forms:
    
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    $Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3$
60. $Fe^{3+}$ ions react with potassium thiocyanate to form a blood-red solution:
    
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    $Fe^{3+} + SCN^- \rightarrow [Fe(SCN)]^{2+}$
    
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Summary

61. The first transition series consists of elements with unique properties due to partially filled $d$-orbitals.
62. They show variable oxidation states, form complex ions, and are coloured.
63. Transition metals, such as iron, play critical roles as catalysts and industrial metals.
64. The extraction of iron involves reducing its oxides in a blast furnace.
65. Steel, an alloy of iron, has superior properties compared to pure iron.
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Jamb chemistry Key points on copper and alloy

Copper (General Information)

1. Copper is a reddish-brown, soft, malleable, and ductile metal.
2. It is a good conductor of heat and electricity.
3. Copper has the chemical symbol Cu and atomic number 29.
4. It is found in nature in the form of ores and native copper.
5. Copper does not corrode easily; it forms a green patina (copper carbonate) when exposed to air for long periods.
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Extraction of Copper from Sulphide Ores

6. The most common copper sulphide ore is chalcopyrite (CuFeS₂).
7. The extraction process involves froth flotation, roasting, smelting, and electrolytic refining.
8. In roasting, sulphide ores are heated in air to convert them into oxides and release sulfur dioxide gas.
   • 2CuFeS₂ + O₂ → Cu₂S + 2FeO + SO₂.
9. The copper(I) sulphide (Cu₂S) is then smelted to form impure copper (blister copper).
10. Impure copper is purified using electrolysis, where pure copper is deposited on the cathode.
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Extraction of Copper from Oxide Ores

11. Oxide ores include cuprite (Cu₂O).
12. Oxide ores are reduced by heating with carbon (coke) in a furnace:
    • Cu₂O + C → 2Cu + CO.
13. Alternatively, oxide ores can undergo leaching with sulfuric acid to form copper(II) sulfate, followed by electrowinning.
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Properties of Copper

14. Copper is an excellent conductor of electricity, second only to silver.
15. It is malleable (can be beaten into thin sheets) and ductile (can be drawn into wires).
16. Copper is resistant to corrosion, making it durable for long-term use.
17. It has a melting point of 1085°C and a boiling point of 2562°C.
18. Copper is non-magnetic and has antimicrobial properties.
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Uses of Copper

19. Copper is used in electrical wiring due to its excellent conductivity.
20. It is used in making pipes and fittings in plumbing.
21. Copper is used in coinage for making coins.
22. It is a key material in heat exchangers and cooling systems.
23. Copper is used in cookware because of its good heat conductivity.
24. It is essential in the production of alloys like brass and bronze.
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Preparation and Uses of Copper(II) Tetraoxosulphate (VI)

25. Copper(II) tetraoxosulphate(VI) or copper(II) sulfate (CuSO₄) is prepared by reacting copper oxide or copper carbonate with sulfuric acid:
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    • CuO + H₂SO₄ → CuSO₄ + H₂O.
26. It is a blue crystalline solid (pentahydrate form: CuSO₄·5H₂O).
27. Copper sulfate is used as an agricultural fungicide.
28. It is used in electroplating copper onto surfaces.
29. It is used in school laboratories for chemical tests and experiments.
30. Copper sulfate is used in the dyeing and printing of textiles.
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Test for Cupric cation

31. Add aqueous sodium hydroxide (NaOH) to a solution containing Cu²⁺ ions. A blue precipitate of copper(II) hydroxide is formed:
    • Cu²⁺ + 2OH⁻ → Cu(OH)₂ (blue ppt).
32. Add aqueous ammonia dropwise to a solution with Cu²⁺. Initially, a blue precipitate forms, which dissolves in excess ammonia to give a deep blue solution.
33. Copper compounds also exhibit a characteristic blue-green flame test.
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Alloy and Its Importance

34. An alloy is a mixture of two or more metals or a metal with a non-metal.
35. Alloys improve properties like strength, durability, resistance to corrosion, and hardness compared to pure metals.
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Steel and Stainless Steel

36. Steel is an alloy of iron and carbon; it is stronger and harder than pure iron.
37. Stainless steel contains iron, chromium, and nickel, making it resistant to rust and corrosion.
38. Stainless steel is used in cutlery, medical instruments, and construction.
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Brass

39. Brass is an alloy of copper and zinc.
40. It is resistant to corrosion and has a golden color, making it suitable for decorative items.
41. Brass is used in musical instruments, plumbing fittings, and locks.
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Bronze

42. Bronze is an alloy of copper and tin.
43. It is harder and more corrosion-resistant than copper.
44. Bronze is used for statues, coins, and machine parts.
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-Metal

45. Type-metal is a lead-based alloy containing antimony and tin.
46. It is used in printing presses because it hardens quickly and retains fine details.
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Duralumin

47. Duralumin is an alloy of aluminum, copper, manganese, and magnesium.
48. It is lightweight and strong, making it suitable for aircraft bodies and automobile parts.
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    ###Soft Solder**
49. Soft solder is an alloy of lead and tin.
50. It is used for joining metals in electronics and plumbing because it has a low melting point.
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Permalloy

51. Permalloy is an alloy of nickel and iron.
52. It is highly magnetic and used in transformers and magnetic shielding.
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Comparing Alloys to Pure Metals

53. Alloys are stronger and harder than pure metals.
54. Alloys often have improved corrosion resistance.
55. Alloys can be tailored for specific applications, e.g., stainless steel for rust resistance.
56. Pure metals like copper and iron are often too soft or reactive for practical use.
57. Adding elements in alloys enhances properties like ductility, toughness, or conductivity.

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