Lesson Notes By Weeks and Term v3 - Senior Secondary 2
Electrolysis
Download the Lessonotes Mobile Nigeria 2025 app for faster lesson access on Android and iPhone.
Electrolysis
Download the Lessonotes Mobile Nigeria 2025 app for faster lesson access on Android and iPhone.
Subject: Chemistry
Class: Senior Secondary 2
Term: 1st Term
Week: 6
Theme: Chemistry And Industry
This page supports the lesson note with a companion video and a short classroom-ready summary.
For class groups and homework, share this lesson page so learners also get the summary, objectives, and full lesson context.
Explain the quantitativeaspects of electrolysis Define electrolytes(strong, weak,fused/molten, nonelectrolytes) electrolytic and electrochemical cells Differentiate betweenstrong and weakelectrolytes Illustrate the electrolysisof acidified water, copper IIsulphates and brines; identify factors affectingthe discharge of ions during Electrolysis Construct the electrolyticand electrochemical cells State Faraday's laws of Electrolysis Calculate amount of substances librated or deposited at electrodesduring electrolysis Explain the uses of electrolysis in extraction and purification of metals
anode. The solution around the cathode becomes alkaline due to the formation of NaOH. This process produces chlorine, hydrogen, and sodium hydroxide, all important industrial chemicals. 2.
8. Faraday's Laws of Electrolysis 2.8.
1. Faraday's First Law: "The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte." Mathematically: m ∝ Q Since Q = It (Quantity of electricity = Current × time), then m ∝ It. So, m = ZIt Where: m = mass of substance deposited/liberated (in grams) Q = quantity of electricity (in Coulombs, C) I = current (in Amperes, A) t = time (in seconds, s) Z = Electrochemical Equivalent (mass of substance deposited by 1 Coulomb of electricity, in g/C). 2.8.
2. Faraday's Second Law: "When the same quantity of electricity is passed through different electrolytes connected in series, the masses of the substances deposited or liberated at the respective electrodes are directly proportional to their chemical equivalents (or equivalent masses)." Mathematically: m1/E1 = m2/E2 = m3/E3 Where: m1, m2, m3 = masses of substances deposited. E1, E2, E3 = chemical equivalents (equivalent masses) of the substances. Equivalent mass (E) = Molar mass (M) / Valency (n) Valency (n) refers to the number of electrons transferred per ion (charge of the ion). So, E = M/n 2.8.
3. Relationship between Faraday's Laws and Fundamental Constants: One Faraday (1 F) is the quantity of electricity required to deposit or liberate one mole of a monovalent substance (i.e., one mole of electrons). 1 F = 96,500 Coulombs (approx.) This means 96,500 C of electricity will discharge 1 mole of a univalent ion (e.g., Na+), 2 moles of electrons for a divalent ion (e.g., Cu2+), and so on.
General formula: Q = nF Where: Q = Quantity of electricity in Coulombs n = number of moles of electrons (also represents valency x moles of substance) F = Faraday's constant (96,500 C/mol e−) Combining this with m = ZIt and Molar mass (M) = E × n, we get: m = (M / nF)
It Where: M is the molar mass of the substance. 2.
9. Uses of Electrolysis
1. Extraction of Metals: Very reactive metals like Aluminium, Sodium, Magnesium, and Calcium are extracted from their fused/molten ores by electrolysis.
Example: Extraction of Aluminium from molten alumina (Al2O3) dissolved in cryolite (Na3AlF6).
Cathode: Al3+ + 3e− → Al(l)
Anode: 2O2− → O2(g) + 4e− (Oxygen reacts with carbon anodes to form CO/CO2) This is crucial for Nigerian students to understand, given the relevance of aluminum in various industries (packaging, construction, vehicle manufacturing).
2. Purification of Metals (Refining): Impure metals are refined by making them the anode in an electrolytic cell, and a pure sample of the same metal is made the cathode.
Example: Purification of copper.
Anode (impure Cu): Cu(s) → Cu2+(aq) + 2e− (and more reactive impurities like Zn, Fe also oxidise)
Cathode (pure Cu): Cu2+(aq) + 2e− → Cu(s) Less reactive impurities (like Ag, Au) fall to the bottom as anode sludge.
3. Electroplating: Coating a cheaper metal with a thin layer of a more expensive or corrosion-resistant metal (e.g., silver, gold, chromium, nickel). The article to be plated is made the cathode. The plating metal is made the anode (or an inert anode is used with a solution of the plating metal's salt).
Examples: Chrome plating for car parts, silver plating for cutlery, gold plating for jewellery. This is a common practice in local craft and repair shops.
4. Manufacture of Chemicals: Chlor-alkali process: Electrolysis of brine to produce chlorine gas, hydrogen gas, and sodium hydroxide.
Production of hydrogen and oxygen: Electrolysis of acidified water. * Production of caustic soda (NaOH): Widely used in soap and detergent industries in Nigeria. 2.
1. Introduction to Electrolysis Electrolysis is the process of using electrical energy to drive non-spontaneous chemical reactions. It involves the decomposition of an electrolyte by passing an electric current through it. This process converts electrical energy into chemical energy. 2.
2. Conductors and Non-conductors Conductors: Materials that allow electricity to pass through them.
Metallic Conductors: Conduct electricity through the movement of free electrons (e.g., copper wires used in household wiring, iron, aluminium). No chemical change occurs.
Electrolytic Conductors (Electrolytes): Conduct electricity through the movement of ions (charged particles). Chemical changes occur at the electrodes.
Non-conductors (Insulators): Materials that do not allow electricity to pass through them (e.g., rubber, plastic, pure water). 2.
3. Electrolytes Electrolytes are substances that produce ions when dissolved in a solvent (usually water) or when melted (fused), and thereby conduct electricity.
Strong Electrolytes: Substances that dissociate completely or almost completely into ions in solution, resulting in high conductivity.
Examples: Strong acids (HCl, H2SO4, HNO3), strong bases (NaOH, KOH), most soluble salts (NaCl, CuSO4).
Weak Electrolytes: Substances that dissociate only partially into ions in solution, resulting in low conductivity.
Examples: Weak acids (CH3COOH, H2CO3), weak bases (NH4OH), sparingly soluble salts.
Fused/Molten Electrolytes: Ionic compounds that become liquid at high temperatures, allowing their ions to move freely and conduct electricity (e.g., molten NaCl, molten Al2O3).
Non-electrolytes: Substances that do not dissociate into ions in solution or when molten, and therefore do not conduct electricity (e.g., sugar, ethanol, urea, pure water). 2.
4. Electrolytic and Electrochemical Cells Electrolytic Cell: Converts electrical energy into chemical energy. Requires an external power source (e.g., battery). Contains two electrodes (anode and cathode) immersed in an electrolyte.
Anode: Positive electrode, where oxidation occurs (anions lose electrons).
Cathode: Negative electrode, where reduction occurs (cations gain electrons). Non-spontaneous reactions occur.
Diagram: [Teacher should draw a simple diagram showing a beaker, two electrodes connected to a battery, and an electrolyte solution]. Electrochemical Cell (Voltaic/Galvanic Cell): Converts chemical energy into electrical energy. Generates electricity from spontaneous redox reactions. Consists of two half-cells, each with an electrode immersed in an electrolyte, connected by a salt bridge and an external wire.
Anode: Negative electrode, where oxidation occurs.
Cathode: Positive electrode, where reduction occurs. Spontaneous reactions occur.
Diagram: [Teacher should draw a simple diagram showing two beakers, electrodes, electrolytes, a salt bridge, and a voltmeter connecting the electrodes]. 2.
5. Mechanism of Electrolysis When an electric current is passed through an electrolyte:
1. Ions Migrate: Cations (positive ions) move towards the cathode (negative electrode), and anions (negative ions) move towards the anode (positive electrode).
2. Discharge at Electrodes: At the cathode, cations gain electrons (reduction) and are discharged. At the anode, anions lose electrons (oxidation) and are discharged. Neutral atoms or molecules are formed. 2.
6. Factors Affecting the Discharge of Ions When more than one type of cation or anion is present, selective discharge occurs based on:
1. Position in the Electrochemical Series (Reactivity Series): At Cathode (Reduction): The ion lower (less reactive) in the electrochemical series is preferentially discharged. Order of discharge (decreasing ease of discharge): Ag+ > Cu2+ > H+ > Pb2+ > Zn2+ > Al3+ > Mg2+ > Na+ > K+. This means Ag+ is more easily reduced than H+, which is more easily reduced than Na+.
At Anode (Oxidation): The ion higher (more reactive) in the electrochemical series is preferentially discharged (if the electrode is inert). For non-metallic ions, the order of discharge is based on ease of oxidation. Order of discharge (decreasing ease of discharge/oxidation): Halides (I− > Br− > Cl−) > OH− > SO42−, NO3−. This means I− is more easily oxidised than Br−, which is more easily oxidised than Cl−, and all are more easily oxidised than OH−. Sulphate and nitrate ions are rarely discharged, instead, OH− from water is discharged.
2. Concentration of Ions: If two ions have similar positions in the electrochemical series, the ion with a higher concentration may be preferentially discharged.
Example:* In the electrolysis of concentrated NaCl solution, oxidation. Order of discharge (decreasing ease of discharge/oxidation): Halides (I− > Br− > Cl−) > OH− > SO42−, NO3−. This means I− is more easily oxidised than Br−, which is more easily oxidised than Cl−, and all are more easily oxidised than OH−. Sulphate and nitrate ions are rarely discharged, instead, OH− from water is discharged.
2. Concentration of Ions: If two ions have similar positions in the electrochemical series, the ion with a higher concentration may be preferentially discharged.
Example: In the electrolysis of concentrated NaCl solution, Cl− ions are discharged at the anode preferentially over OH− ions, despite OH− being slightly higher in the reactivity series for anions. In dilute NaCl, OH− is discharged.
3. Nature of the Electrode: Inert Electrodes: (e.g., platinum, graphite) do not participate in the reaction; they only provide a surface for electron transfer.
Active Electrodes: (e.g., copper, silver) participate in the reaction by themselves dissolving or reacting. 2.
7. Illustrations of Electrolysis 2.7.
1. Electrolysis of Acidified Water (using inert electrodes like platinum)
Electrolyte: Water is a weak electrolyte. A small amount of acid (e.g., H2SO4) is added to increase its conductivity by providing more H+ and SO42− ions.
Ions Present: H+, OH− (from water), H+, SO42− (from acid).
At Cathode (Negative electrode): H+ ions are discharged. 2H+(aq) + 2e− → H2(g)
At Anode (Positive electrode): OH− ions are discharged (as SO42− is much harder to oxidise). 4OH−(aq) → O2(g) + 2H2O(l) + 4e− Overall Reaction: 2H2O(l) → 2H2(g) + O2(g)
Observations: Hydrogen gas is collected at the cathode (twice the volume of oxygen). Oxygen gas is collected at the anode. The electrolyte remains acidic. 2.7.
2. Electrolysis of Aqueous Copper(II)
Sulphate Solution Case A: Using inert electrodes (e.g., platinum or graphite)
Ions Present: Cu2+, SO42− (from CuSO4), H+, OH− (from water).
At Cathode: Cu2+ is lower than H+ in the electrochemical series, so Cu2+ is discharged. Cu2+(aq) + 2e− → Cu(s)
At Anode: OH− is preferentially discharged over SO42−. 4OH−(aq) → O2(g) + 2H2O(l) + 4e− Observations: A pinkish-brown deposit of copper forms on the cathode. Oxygen gas is evolved at the anode. The blue colour of the solution fades as Cu2+ ions are removed, and the solution becomes acidic due to the accumulation of H+ ions.
Case B: Using active copper electrodes Ions Present: Cu2+, SO42− (from CuSO4), H+, OH− (from water).
At Cathode: Cu2+ is discharged, and copper metal is deposited. Cu2+(aq) + 2e− → Cu(s)
At Anode: The copper anode itself oxidises (dissolves) preferentially over OH− or SO42− ions. Cu(s) → Cu2+(aq) + 2e− Observations: A pinkish-brown deposit of copper forms on the cathode. The anode loses mass as copper dissolves. The concentration and blue colour of the solution remain relatively constant as Cu2+ ions are deposited at the cathode and replenished from the anode. This is the principle of refining crude copper. 2.7.
3. Electrolysis of Brine (Concentrated Aqueous Sodium Chloride)
Ions Present: Na+, Cl− (from NaCl), H+, OH− (from water).
At Cathode: H+ is preferentially discharged over Na+ (H+ is lower in the series). 2H+(aq) + 2e− → H2(g) Alternatively, 2H2O(l) + 2e− → H2(g) + 2OH−(aq)
At Anode: Due to high concentration, Cl− is preferentially discharged over OH− (even though OH− is slightly higher in the series in terms of ease of oxidation). 2Cl−(aq) → Cl2(g) + 2e− Overall Reaction: 2NaCl(aq) + 2H2O(l) → 2NaOH(aq) + H2(g) + Cl2(g)
Observations: Hydrogen gas is evolved at the cathode. Chlorine gas is evolved at the anode. The solution around the cathode becomes alkaline due to the formation of NaOH. This process produces chlorine, hydrogen, and sodium hydroxide, all important industrial chemicals. 2.
8. Faraday's Laws of Electrolysis 2.8.
1. Faraday's First Law: "The mass of a substance deposited or liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte." Mathematically: m ∝ Q Since Q = It (Quantity of electricity = Current × time), then m ∝ It. So, m = ZIt Where: m = mass of substance deposited/liberated (in 3.
1. Introduction (10 minutes)
Teacher Activity: Begin by reviewing conductors and insulators. Ask students to identify examples of materials that allow electricity to flow and those that do not. Introduce the concept of chemical reactions driven by electricity.
Student Activity: Students brainstorm examples of conductors and insulators. Answer questions on the difference between metallic conduction and conduction through solutions. 3.
2. Definition and Types of Electrolytes (15 minutes)
Teacher Activity: Present the definitions of electrolytes, non-electrolytes, strong electrolytes, weak electrolytes, and fused/molten electrolytes. Provide examples for each and use a simple conductivity apparatus (battery, bulb, electrodes) to demonstrate the conductivity of distilled water, tap water, salt solution, sugar solution, dilute acid, and dilute base.
Student Activity: Students record definitions and examples. Observe conductivity tests and classify substances based on their conductivity. Discuss why some solutions conduct better than others. 3.
3. Electrolytic vs. Electrochemical Cells (15 minutes)
Teacher Activity: Present diagrams and explain the components and functions of electrolytic and electrochemical cells. Emphasise the energy conversion (electrical to chemical vs. chemical to electrical), spontaneity, and the polarity of electrodes.
Student Activity: Students draw and label diagrams of both cell types. Complete a Venn diagram or a comparison table to highlight the similarities and differences. 3.
4. Mechanism and Factors Affecting Discharge (20 minutes)
Teacher Activity: Explain the migration of ions and discharge at electrodes. Introduce the electrochemical series (reactivity series) for cations and anions. Explain how concentration and electrode nature influence selective discharge.
Student Activity: Students write down the general reactions at anode and cathode. They will use the electrochemical series to predict products of hypothetical electrolysis scenarios. 3.
5. Practical Illustrations of Electrolysis (30 minutes)
Teacher Activity: Demonstration 1: Electrolysis of Acidified Water. Set up a Hoffman voltameter or a simpler apparatus with test tubes inverted over platinum electrodes in dilute H2SO
4. Observe gas collection. (If apparatus is unavailable, use visual aids/videos).
Demonstration 2: Electrolysis of Copper(II) Sulphate. Use inert graphite electrodes and then, if possible, copper electrodes. Observe colour changes, deposits, and gas evolution.
Discussion: Guide students to write down ionic equations for each electrode reaction and the overall reaction for all demonstrations. Explain the observations and link them to the factors affecting discharge. For brine, discuss the industrial relevance of products (NaOH, Cl2, H2).
Student Activity: Students actively observe the demonstrations. Record observations, predict products, and write balanced chemical equations for electrode reactions and overall processes. Engage in discussion about real-world uses of the products (e.g., chlorine in water treatment, caustic soda in soap making). 3.
6. Faraday's Laws of Electrolysis (30 minutes)
Teacher Activity: Introduce Faraday's First Law with the formula m = ZIt. Define Z and its units. Introduce Faraday's Second Law and the concept of equivalent mass (E = M/n). Explain the constant F = 96,500 C/mol. Show how to combine these concepts to solve quantitative problems using clear, step-by-step calculations.
Student Activity: Students record the laws and formulas. Practice rearranging formulas and identifying variables. Work through simple calculation examples provided by the teacher. 3.
7. Uses of Electrolysis (15 minutes)
Teacher Activity: Discuss the industrial applications of electrolysis: extraction and purification of metals (e.g., aluminium from bauxite, copper refining), electroplating, and manufacture of chemicals (e.g., chlor-alkali process, production of H2 and O2). Relate these uses to Nigerian industries or daily life examples.
Student Activity: Students take notes on the applications. They will brainstorm examples of electroplated items they use or see in Nigeria (e.g., jewellery, car parts, roofing sheets).
Water Purification and Treatment: Electrolysis, particularly using electrocoagulation or electrochlorination, can be used to treat water by removing suspended particles, heavy metals, and microorganisms. In Nigeria, where access to clean water can be a challenge in some rural and urban areas, understanding these electrochemical methods is crucial for local community water projects and could inform sustainable solutions for portable water provision.
Metal Industries and Manufacturing: The extraction of aluminium from its ore (bauxite) and the refining of copper (for electrical cables, which are critical for power distribution in Nigeria) are massive industrial processes relying on electrolysis. Students can relate this to the importance of metal industries for infrastructure development, manufacturing (e.g., aluminium pots, roofing sheets), and job creation in Nigeria. Electroplating is also prevalent in local workshops for protecting metals from corrosion or for decorative purposes (e.g., gold-plated jewellery, chrome-plated vehicle parts).
Chemical Production for Local Industries: The chlor-alkali process, which uses electrolysis of brine, produces sodium hydroxide (caustic soda), chlorine gas, and hydrogen gas. Sodium hydroxide is a key ingredient in local soap and detergent manufacturing industries in Nigeria. Chlorine is used for water disinfection and in the plastics industry. This directly links the abstract chemical process to tangible products encountered daily and economic activities within the country.