ICSE Class 10 Chemistry Redox Reactions and Electrolysis Notes
Mastering ICSE Class 10 Chemistry Redox Reactions and Electrolysis Notes is essential for scoring top marks in your board examinations. Chemical reactions involving electron transfer form the backbone of physical and inorganic chemistry. From understanding how metals corrode to exploring the industrial extraction of aluminum and electroplating delicate jewelry, redox chemistry and electrolysis govern critical laboratory and industrial transformations.
This comprehensive revision guide covers core definitions, electronic interpretations of oxidation and reduction, preferential discharge theory, complete electrode mechanisms, industrial applications, and official board exam questions with detailed model solutions.
Fundamental Concepts of Redox Reactions
A redox (reduction-oxidation) reaction is a chemical process in which oxidation and reduction occur simultaneously. One substance cannot lose electrons or oxygen unless another substance is present to accept them.
1. Classical Concept (Oxygen and Hydrogen Transfer)
- Oxidation: The addition of oxygen or the removal of hydrogen from a substance.
- Reduction: The addition of hydrogen or the removal of oxygen from a substance.
Example: In the reaction CuO + H2 → Cu + H2O, copper(II) oxide loses oxygen (reduced to Cu), while hydrogen gains oxygen (oxidized to H2O).
2. Modern Electronic Concept (Electron Transfer)
The electronic concept provides a universal definition applicable to all ionic and covalent processes:
- Oxidation (OIL - Oxidation Is Loss): A process in which an atom, ion, or molecule loses one or more electrons, resulting in an increase in positive valency or oxidation state.
Fe2+ → Fe3+ + e−(Oxidation of Ferrous to Ferric ion) - Reduction (RIG - Reduction Is Gain): A process in which an atom, ion, or molecule gains one or more electrons, resulting in a decrease in positive valency or an increase in negative valency.
Cu2+ + 2e− → Cu(Reduction of Cupric ion to Copper metal)
3. Oxidizing and Reducing Agents
- Oxidizing Agent (Oxidant): A chemical species that oxidizes another substance by accepting electrons (it gets reduced in the process). Examples: Concentrated HNO3, acidified KMnO4, Cl2, O2, and H2O2.
- Reducing Agent (Reductant): A chemical species that reduces another substance by donating electrons (it gets oxidized in the process). Examples: Active metals (Na, Zn, Fe), H2, CO, SO2, and H2S.
Understanding Electrolysis as an Applied Redox Process
Electrolysis is the chemical decomposition of an electrolyte in its molten state or aqueous solution by the passage of a direct electric current (DC). It represents a non-spontaneous redox reaction driven by external electrical energy.
Electrolytes vs. Non-Electrolytes
- Strong Electrolytes: Substances that dissociate almost completely into ions in aqueous solution or molten state, conducting large currents (e.g., dilute H2SO4, HCl, NaOH, molten NaCl).
- Weak Electrolytes: Substances that dissociate only partially into ions in solution, consisting of both ions and unionized molecules (e.g., CH3COOH, NH4OH, H2CO3).
- Non-Electrolytes: Liquid compounds that do not form ions and do not conduct electricity (e.g., pure distilled water, liquid carbon tetrachloride CCl4, alcohol, glucose solution).
The Electrolytic Cell Architecture
An electrolytic cell converts electrical energy into chemical energy through two specific electrodes:
- Anode (Positive Electrode): Connected to the positive terminal of the battery. Anions (negative ions) migrate to the anode and lose electrons (Oxidation occurs at the Anode — Mnemonic: AN OX).
- Cathode (Negative Electrode): Connected to the negative terminal of the battery. Cations (positive ions) migrate to the cathode and gain electrons (Reduction occurs at the Cathode — Mnemonic: RED CAT).
Selective Discharge of Ions (Preferential Discharge Theory)
When an electrolyte contains more than one type of cation or anion, all ions migrate toward their respective electrodes, but only one type of ion discharges at a given electrode at a time. The preferential discharge depends on three major factors:
1. Position of the Ion in the Electrochemical Series
Ions placed lower in the electrochemical series require less energy to discharge and are therefore preferentially discharged over ions placed higher up.
- Cations (Ease of discharge increases downwards):
K+ < Ca2+ < Na+ < Mg2+ < Al3+ < Zn2+ < Fe2+ < Pb2+ < H+ < Cu2+ < Ag+ < Au3+ - Anions (Ease of discharge increases downwards):
SO42− < NO3− < Cl− < Br− < I− < OH−
2. Concentration of Ions in Solution
A higher concentration of a specific ion can overcome a slight difference in electrochemical positioning (e.g., in concentrated NaCl/brine solution, Cl− is discharged at the anode in preference to OH− due to high concentration).
3. Nature of the Electrodes (Inert vs. Active)
Inert electrodes (such as Platinum or Graphite) do not take part in the chemical reaction. Active electrodes (such as Copper, Nickel, or Silver) dissolve at the anode during electrolysis, participating directly in the redox cycle.
Detailed Examples of Electrolysis
1. Electrolysis of Molten Lead Bromide (PbBr2)
- Electrolyte: Molten PbBr2 (maintained at ~380°C in a silica crucible).
- Electrodes: Graphite rods (inert to reactive bromine vapors).
- Dissociation:
PbBr2 ⇔ Pb2+ + 2Br− - Reaction at Cathode (Reduction):
Pb2+ + 2e− → Pb(l)→ Silvery grey metal deposits at the cathode. - Reaction at Anode (Oxidation):
2Br− − 2e− → Br2(g)→ Reddish-brown bromine fumes evolve at the anode.
2. Electrolysis of Acidified Water (Hoffmann's Voltameter)
- Electrolyte: Distilled water acidified with dilute H2SO4 (since pure water is a non-electrolyte).
- Electrodes: Platinum foils.
- Ions Present: H+, OH− (from H2O) and H+, SO42− (from H2SO4).
- At Cathode:
2H+ + 2e− → H2(g)(2 volumes of Hydrogen gas collected). - At Anode: OH− discharges preferentially over SO42−:
4OH− − 4e− → 2H2O + O2(g)(1 volume of Oxygen gas collected). - Volume Ratio: Ratio of H2 : O2 produced is strictly 2 : 1 by volume.
3. Electrolysis of Aqueous Copper(II) Sulphate
- Case A: Using Platinum (Inert) Electrodes:
- Cathode:
Cu2+ + 2e− → Cu(s)(Pinkish-brown copper deposits). - Anode:
4OH− − 4e− → 2H2O + O2(g)(Oxygen gas evolved). - Observation: The blue color of the electrolyte gradually fades as Cu2+ ions are consumed without replenishment.
- Cathode:
- Case B: Using Copper (Active) Electrodes:
- Cathode:
Cu2+ + 2e− → Cu(s)(Cathode increases in mass). - Anode:
Cu(s) − 2e− → Cu2+(aq)(Anode dissolves and decreases in mass). - Observation: The blue color of CuSO4 remains unchanged because the rate of loss of Cu2+ at the cathode equals the rate of production of Cu2+ at the anode.
- Cathode:
Industrial Applications of Electrolysis
- Electroplating: Coating a superior metal (e.g., Ni, Ag, Au, Cr) onto a base metal (e.g., iron, brass) to prevent corrosion and improve appearance.
- The article to be electroplated is always made the Cathode.
- The pure plating metal is made the Anode.
- The electrolyte must contain a soluble salt of the plating metal.
- Electro-refining of Metals: Purifying crude blister copper. Impure copper acts as the anode, thin pure copper sheet acts as the cathode, and acidified CuSO4 is the electrolyte. Pure copper transfers to the cathode while impurities settle as anode sludge/mud.
- Electrometallurgy: Extraction of highly electropositive metals (K, Na, Ca, Mg, Al) from their fused halides or oxides (e.g., extraction of Aluminium via the Hall-Héroult process).
Important Board Exam Questions with Answers
The following questions represent core conceptual and equation-based problems frequently asked in Class 10 board examinations:
Question 1: Define exothermic reactions. Give one example.
Answer: An exothermic reaction is a chemical reaction that releases heat energy into the surroundings along with the formation of products.
Example: Combustion of natural gas (methane):
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + Heat
Question 2: Why is respiration considered an exothermic reaction?
Answer: During respiration, the carbohydrates digested into glucose combine with oxygen in the cells of our body to produce carbon dioxide, water, and energy in the form of ATP and heat:
C6H12O6(aq) + 6O2(g) → 6CO2(g) + 6H2O(l) + Energy
Because energy is released during the breakdown of glucose molecules, respiration is classified as an exothermic process.
Question 3: What happens when zinc granules are treated with dilute H2SO4?
Answer: When zinc granules react with dilute sulphuric acid, single displacement occurs yielding zinc sulphate and liberating hydrogen gas with brisk effervescence:
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)↑
Observation: Bubbles of colourless, odourless H2 gas evolve, which burn with a characteristic 'pop' sound when tested with a burning splinter.
Question 4: Define rancidity. How can it be prevented?
Answer: Rancidity is the slow aerial oxidation of fats and oils present in food items, resulting in the formation of volatile aldehydes and carboxylic acids that cause an unpleasant smell and foul taste.
Methods of Prevention:
- Adding synthetic antioxidants such as BHA (Butylated Hydroxyanisole) and BHT (Butylated Hydroxytoluene).
- Flushing food packaging bags with an inert gas like Nitrogen (N2) to displace oxygen.
- Storing food in airtight containers away from direct sunlight.
- Refrigeration of perishable fatty foods to slow down oxidation kinetics.
Question 5: What is a displacement reaction? Give a chemical equation to illustrate.
Answer: A displacement reaction is a chemical process where a more reactive element displaces a less reactive element from its aqueous salt solution.
Equation:
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Iron (Fe) is placed higher in the activity series than copper (Cu); hence, iron displaces copper, turning the blue copper sulphate solution into light pale-green ferrous sulphate while reddish-brown copper deposits on the iron nail.
Question 6: Balance the equation: Fe + H2O → Fe3O4 + H2
Answer:
Step 1: Balance Fe atoms by placing coefficient 3 before Fe: 3Fe + H2O → Fe3O4 + H2
Step 2: Balance Oxygen atoms by placing coefficient 4 before H2O: 3Fe + 4H2O → Fe3O4 + H2
Step 3: Balance Hydrogen atoms by placing coefficient 4 before H2: 3Fe + 4H2O → Fe3O4 + 4H2
Final Balanced Equation:
3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)
Question 7: Why do silver chloride and silver bromide turn grey on exposure to sunlight?
Answer: Silver chloride (AgCl) and silver bromide (AgBr) undergo photochemical decomposition in the presence of light photons (hν). White silver chloride and pale-yellow silver bromide break down into elemental silver metal and halogen gases:
2AgCl(s) → 2Ag(s) [grey] + Cl2(g)2AgBr(s) → 2Ag(s) [grey] + Br2(g)
The formation of finely divided, metallic grey silver particles causes the visible colour transition.
Question 8: Identify the type of chemical reaction in each case and write balanced equations: (a) Decomposition of calcium carbonate on heating (b) Reaction of sodium with water
Answer:
(a) Decomposition of calcium carbonate on heating:
Type of reaction: Thermal Decomposition Reaction
Balanced Equation: CaCO3(s) → CaO(s) + CO2(g)↑
(b) Reaction of sodium with water:
Type of reaction: Exothermic Single Displacement Reaction (Redox)
Balanced Equation: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)↑ + Heat
How to Prepare for Redox Reactions and Electrolysis
- Master Half-Cell Equations: Always separate ionic redox reactions into oxidation and reduction half-reactions. Check that both atoms and net electrical charges balance on both sides.
- Memorize the Electrochemical Discharge Hierarchy: Practice predicting which ion discharges first in mixed electrolytes (e.g., dilute aqueous NaCl vs concentrated brine).
- Understand the Dual Role of Electrodes: Differentiate clearly between inert electrodes (Pt, graphite) where the electrode only transfers electrons, and active electrodes (Cu, Ag, Ni) where the anode actively oxidizes into solution.
- Map Out Practical Applications: Create schematic diagrams for electroplating an iron spoon with silver, refining blister copper, and the extraction of molten lead bromide. Note the electrolyte, cathode, anode, and observation for each setup.
Where to Practice More
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