Electrolysis | Jamb Chemistry

Jamb chemistry key points on Electrolytes and non-electrolytes; Faraday’s laws of electrolysis etc

Electrolytes and Non-Electrolytes

1. Electrolytes conduct electricity in molten or aqueous states.
2. Examples include acids (HCl), bases (NaOH), and salts (NaCl).
3. Electrolytes dissociate into ions in solution.
4. Non-electrolytes do not conduct electricity.
5. Examples of non-electrolytes are glucose and urea.
6. Strong electrolytes fully dissociate into ions.
7. Weak electrolytes partially dissociate into ions.
8. Non-electrolytes dissolve as neutral molecules.
9. Electrolyte strength determines conductivity.
10. Conductivity tests differentiate electrolytes from non-electrolytes.
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Faraday’s Laws of Electrolysis

11. First Law: The mass of a substance deposited is proportional to the charge passed.
12. Second Law: The masses of substances deposited are proportional to their equivalent weights.
13. Mathematically: $m = \frac{Q}{F} \times \frac{M}{z}$.
14. $Q$ is charge in coulombs, $F$ is Faraday’s constant, $M$ is molar mass, $z$ is valency.
15. 1 Faraday equals $96,485C/mol$.
16. The number of moles of electrons is $\frac{Q}{F}$.
17. Faraday’s laws apply to industrial electrolysis.
18. Electrochemical equivalents guide production rates.
19. The laws predict deposition amounts in electroplating.
20. They are essential for accurate electrolysis calculations.
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Electrolysis of Various Solutions

Dilute H₂SO₄

21. At the cathode: $2H^+ + 2e^- \rightarrow H_2$.
22. At the anode: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$.
23. Products are hydrogen ($H_2$) and oxygen ($O_2$).
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Aqueous CuSO₄ (Inert Electrodes)

24. Cathode: $Cu^{2+} + 2e^- \rightarrow Cu$.
25. Anode: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$.
26. Products are copper ($Cu$) and oxygen ($O_2$).
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Aqueous CuSO₄ (Copper Electrodes)

27. Anode: $Cu \rightarrow Cu^{2+} + 2e^-$.
28. Cathode: $Cu^{2+} + 2e^- \rightarrow Cu$.
29. Copper is transferred from anode to cathode.
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Aqueous CuCl₂

30. Cathode: $Cu^{2+} + 2e^- \rightarrow Cu$.
31. Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$.
32. Products are copper ($Cu$) and chlorine gas ($Cl_2$).
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Dilute NaCl

33. Cathode: $2H_2O + 2e^- \rightarrow H_2 + 2OH^-$.
34. Anode: $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$.
35. Products are hydrogen ($H_2$) and oxygen ($O_2$).
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Concentrated NaCl

36. Cathode: $2H_2O + 2e^- \rightarrow H_2 + 2OH^-$.
37. Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$.
38. Products are hydrogen ($H_2$) and chlorine ($Cl_2$).
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Fused NaCl

39. Cathode: $Na^+ + e^- \rightarrow Na$.
40. Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$.
41. Products are sodium ($Na$) and chlorine ($Cl_2$).
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Factors Affecting Discharge of Ions

42. Ion concentration influences discharge.
43. Electrode material determines product formation.
44. Lower ions in the electrochemical series discharge preferentially.
45. Overvoltage can affect ion discharge sequence.
46. Presence of complexing agents modifies ion discharge.
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Calculations Based on Faraday

47. Charge is calculated as $Q = I \times t$, where $I$ is current, $t$ is time.
48. Moles of electrons are $\frac{Q}{F}$.
49. Amount of product is proportional to moles of electrons.
50. $m = \frac{Q}{F} \times \frac{M}{z}$ determines deposited mass.
51. For $Cu^{2+}$, $Q = 2 \times F$ per mole of copper.
52. Faraday’s constant ($96,485C/mol$) relates charge to substance.
53. Time for electrolysis is $t = \frac{Q}{I}$.
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Suitable Electrodes for Different Electrolytes

54. Inert electrodes like platinum and graphite are non-reactive.
55. Active electrodes like copper participate in reactions.
56. Graphite is suitable for NaCl and H₂SO₄ electrolysis.
57. Copper electrodes are ideal for CuSO₄ electrolysis.
58. Nickel electrodes are used for electroplating.
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Chemical Reactions at the Electrodes

59. Cathode reactions are reductions (e.g., $Cu^{2+} + 2e^- \rightarrow Cu$).
60. Anode reactions are oxidations (e.g., $2Cl^- \rightarrow Cl_2 + 2e^-$).
61. Hydrogen gas evolves at cathodes in water-based solutions.
62. Oxygen forms at anodes from water oxidation.
63. Chlorine forms at the anode from chloride ions.
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Determination of Products at Electrodes

64. Ion type determines electrode products.
65. The electrochemical series predicts discharge.
66. Concentration affects which ions are discharged.
67. Water may compete with ions for discharge in aqueous solutions.
68. Overvoltage alters predicted products.
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Uses of Electrolysis

Purification of Metals

69. Electrolysis purifies copper in refining processes.
70. Impure copper acts as the anode.
71. Pure copper is deposited at the cathode.
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Production of Elements

72. Electrolysis of NaCl produces sodium and chlorine.
73. Water electrolysis yields hydrogen and oxygen.
74. Alumina electrolysis produces aluminum.
75. Brine electrolysis produces chlorine, hydrogen, and NaOH.
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Electroplating

76. Electrolysis deposits thin metal layers on objects.
77. The object to be plated is the cathode.
78. Electroplating enhances durability and appearance.
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Battery Technology

79. Rechargeable batteries use electrolysis for energy storage.
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Corrosion Protection

80. Cathodic protection uses electrolysis to prevent rust.
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Industrial Applications

81. Hydrogen is produced for fuel cells.
82. Chlorine gas is essential for PVC production.
83. Sodium hydroxide is used in soaps and detergents.
84. Aluminum production uses the Hall-Héroult process.
85. Water purification employs electrolysis.
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Biological and Environmental Relevance

86. Electrolysis purifies water in desalination plants.
87. It plays a role in wastewater treatment.
88. Renewable energy powers eco-friendly electrolysis.
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Experimental Observations

89. Bubble formation at electrodes indicates gas production.
90. Color changes reveal ion movement.
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Miscellaneous

91. Faraday’s laws are fundamental in electrochemistry.
92. Electrolysis confirms the electrolyte’s activity.
93. Conductivity depends on ion presence.
94. Industrial electrolysis enhances resource efficiency.
95. Electrolysis is key in chemical synthesis.
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Predicting and Testing

96. Conductivity tests identify electrolytes.
97. Ion migration depends on electrode polarity.
98. The electrochemical series guides discharge predictions.
99. Electrode material affects reaction type and efficiency.
100. Understanding electrolysis is essential for practical applications in chemistry.
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Jamb chemistry Key points on Electrochemical cells; Corrosion as an electrolytic process

Electrochemical Cells

1. Electrochemical cells convert chemical energy into electrical energy.
2. They consist of two electrodes: the anode and the cathode.
3. Oxidation occurs at the anode, releasing electrons.
4. Reduction occurs at the cathode, accepting electrons.
5. The flow of electrons creates an electric current.
6. A salt bridge maintains charge balance by allowing ion flow between half-cells.
7. A galvanic cell is a type of electrochemical cell that generates electricity spontaneously.
8. In a Daniell cell, zinc is oxidized ($Zn \rightarrow Zn^{2+} + 2e^-$), and copper is reduced ($Cu^{2+} + 2e^- \rightarrow Cu$).
9. The cell potential is the difference in electrode potentials between the two half-cells.
10. Standard conditions for cell operation are 1 M concentration, 25°C, and 1 atm pressure.
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Electrochemical Series

11. The electrochemical series ranks elements by their electrode potentials.
12. Elements at the top (e.g., $K, Ca, Na$) are strong reducing agents.
13. Elements at the bottom (e.g., $Ag, Au$) are strong oxidizing agents.
14. Potassium ($K$) has the most negative electrode potential and is highly reactive.
15. Gold ($Au$) has the most positive electrode potential and is least reactive.
16. Metals above hydrogen can displace $H_2$ gas from acids.
17. Metals below hydrogen (e.g., $Cu, Hg, Ag$) cannot displace $H_2$ gas.
18. The position in the series determines a metal’s reactivity in redox reactions.
19. Zinc reacts with dilute acids, as it is above hydrogen in the series.
20. Copper does not react with dilute acids, as it is below hydrogen.
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Simple Calculations on Half-Cell Reactions and Electrode Potentials**

21. The cell potential ($E_{cell}$) is calculated as:
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    $E_{cell} = E_{cathode} - E_{anode}$
22. A positive $E_{cell}$ indicates a spontaneous reaction.
23. Electrode potential is measured against the standard hydrogen electrode (SHE).
24. For $Zn$ and $Cu$ in a Daniell cell:
    
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    $E_{cell} = E_{Cu^{2+}/Cu} - E_{Zn^{2+}/Zn}$
25. If $E_{cell} > 0$, the cell can generate electricity.
26. Standard electrode potentials help predict redox reaction feasibility.
27. The Nernst equation calculates potentials under non-standard conditions:
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    $E = E^\circ - \frac{RT}{nF} \ln Q$
28. Half-cell reactions are balanced for mass and charge to determine $E_{cell}$.
29. Gibbs free energy is related to cell potential:
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    $\Delta G = -nFE_{cell}$
30. A negative $\Delta G$ indicates a thermodynamically favorable reaction.
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Corrosion as an Electrolytic Process

31. Corrosion is the deterioration of metals due to chemical reactions with the environment.
32. It involves electrochemical reactions where the metal acts as an anode.
33. At the anodic site, metal oxidizes to ions (e.g., $Fe \rightarrow Fe^{2+} + 2e^-$).
34. At the cathodic site, oxygen is reduced ($O_2 + 4H^+ + 4e^- \rightarrow 2H_2O$).
35. Moisture and electrolytes accelerate corrosion.
36. Iron forms rust ($Fe_2O_3 \cdot xH_2O$) due to reaction with oxygen and water.
37. Corrosion weakens metal structures and reduces their lifespan.
38. High humidity, salts, and acids exacerbate corrosion.
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Cathodic Protection of Metals

39. Cathodic protection prevents corrosion by making the metal act as a cathode.
40. Sacrificial anodes (e.g., $Zn, Mg$) corrode in place of the protected metal.
41. The sacrificial anode donates electrons to the protected metal.
42. Impressed current systems use an external power source to supply electrons.
43. Cathodic protection is widely used for pipelines and ship hulls.
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Painting, Electroplating, and Coating with Grease or Oil to Prevent Corrosion

44. Painting creates a barrier, protecting metal from air and moisture.
45. Paint coatings are used in automotive and construction industries.
46. Electroplating coats metals with a thin protective layer (e.g., chrome or zinc).
47. Electroplating improves corrosion resistance and aesthetic appeal.
48. Grease and oil act as temporary barriers against moisture.
49. Lubricants protect moving parts from rusting.
50. These methods enhance metal durability and reduce maintenance costs.
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Applications of Electrolytic Processes

Industrial Applications

51. Electrolysis is used to produce metals like aluminum (Hall-Héroult process).
52. Sodium is obtained via electrolysis of molten NaCl.
53. Electrolysis of brine produces chlorine, hydrogen, and NaOH.
54. Purification of copper involves electrolysis in electrolytic refining.
55. Zinc is purified using electrolytic methods.
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Electroplating

56. Electroplating enhances durability and corrosion resistance.
57. Jewelry is electroplated with gold or silver for aesthetic value.
58. Chromium plating provides a shiny, hard surface.
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Water Treatment

59. Electrolysis purifies water by breaking down contaminants.
60. Desalination plants use electrolytic processes for fresh water production.
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Energy Storage

61. Electrolytic cells are used in rechargeable batteries (e.g., $Li$-ion batteries).
62. Hydrogen fuel cells involve electrolysis to produce $H_2$ for clean energy.
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Electrochemical Sensors

63. Electrochemical cells detect gases like oxygen and carbon dioxide.
64. Glucose sensors for diabetes monitoring use electrochemical principles.
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Biological and Environmental Applications

65. Electrochemical processes help in wastewater treatment.
66. Electrolysis removes heavy metals from industrial effluents.
67. It is used in the electrochemical remediation of polluted soil.
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Experimental and Laboratory Uses

68. Electrolysis demonstrates redox reactions in educational experiments.
69. Electrochemical cells measure thermodynamic properties of reactions.
70. Electrolysis confirms the presence of electrolytes in solutions.
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Corrosion Prevention Beyond Cathodic Protection

71. Alloying metals like stainless steel reduces corrosion susceptibility.
72. Anodization forms a protective oxide layer on aluminum.
73. Galvanizing coats steel with zinc for corrosion resistance.
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Advanced Electrochemical Applications

74. Electrochemical processes synthesize chemicals like hydrogen peroxide.
75. Electrolysis produces industrial gases like oxygen and chlorine.
76. Electrochemical machining shapes metals with precision.
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Miscellaneous

77. Electrochemical series predicts reactivity trends in metals.
78. Understanding half-cell reactions optimizes battery efficiency.
79. Corrosion control saves industries billions annually.
80. Electrochemical applications are vital in sustainable energy solutions.

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