Class 9 Science Chapter 11: Chemical Effects of Electric Current – Important Questions with Complete Solutions

The CBSE Class 9 Science curriculum (2024–25 rationalized syllabus) emphasizes application-based learning alongside conceptual understanding. Chapter 11, Chemical Effects of Electric Current, typically contributes 6–8 marks to the annual board exam. Recent papers show a balanced split: ~40% objective questions (MCQs, fill-in-the-blanks), ~35% short-answer (2–3 marks), and ~25% long-answer (5 marks). Examiners focus on practical scenarios—why silver articles are electroplated, how to test whether a liquid is a conductor, reactions occurring at electrodes during electrolysis. Memorizing definitions is insufficient; students must explain *why* a solution conducts electricity (ion movement), *how* electroplating deposits metal, and *what* products form when specific liquids are electrolysed. The questions in this guide reflect these patterns precisely, ensuring you're ready for term exams and the final board assessment.

**Question 1:** Which of the following liquids is the best conductor of electricity? (a) Distilled water (b) Tap water (c) Kerosene (d) Acetone **Answer:** (b) Tap water **Explanation:** Tap water contains dissolved salts and minerals that ionize, producing free ions. These ions carry electric charge, making tap water conductive. Distilled water has no ions, so it is a poor conductor. Kerosene and acetone are non-polar organic liquids with negligible ionic content. --- **Question 2:** During electrolysis of water, the gas evolved at the anode is: (a) Hydrogen (b) Oxygen (c) Chlorine (d) Nitrogen **Answer:** (b) Oxygen **Explanation:** At the anode (positive electrode), negatively charged hydroxide ions (OH⁻) lose electrons and oxidize to oxygen gas: 4OH⁻ → O₂ + 2H₂O + 4e⁻. At the cathode, hydrogen gas is formed. --- **Question 3:** The process of coating an object with a layer of metal by electrolysis is called: (a) Ionization (b) Electroplating (c) Oxidation (d) Reduction **Answer:** (b) Electroplating **Explanation:** Electroplating uses electrolysis to deposit a thin, even layer of one metal onto another. The object to be plated is the cathode, the coating metal is the anode, and the electrolyte is a salt solution of the coating metal. --- **Question 4:** In electrolysis, which electrode is the positive terminal of the battery? (a) Anode (b) Cathode (c) Both anode and cathode (d) Neither anode nor cathode **Answer:** (a) Anode **Explanation:** The anode is connected to the positive terminal and attracts anions (negative ions). The cathode connects to the negative terminal and attracts cations (positive ions). --- **Question 5:** A liquid that conducts electricity is called a(n): (a) Electrolyte (b) Nonelectrolyte (c) Insulator (d) Semiconductor **Answer:** (a) Electrolyte **Explanation:** An electrolyte is a liquid or solution containing free ions capable of conducting electricity. Examples: dilute acids, bases, and salt solutions.

**Question 1:** Why is distilled water a poor conductor of electricity, while tap water is a good conductor? Explain. **Answer:** Distilled water is pure H₂O with virtually no ions. In the absence of free charge carriers, it cannot conduct electric current. Tap water, however, contains dissolved salts like NaCl, CaSO₄, and minerals. These salts dissociate into ions (Na⁺, Cl⁻, Ca²⁺, SO₄²⁻, etc.). Free ions move toward the electrodes and carry electric charge, making tap water conductive. --- **Question 2:** In the electrolysis of copper sulphate (CuSO₄) solution using copper electrodes, what happens at each electrode? Write the equations. **Answer:** - **At the cathode (−):** Cu²⁺ + 2e⁻ → Cu (copper deposits) - **At the anode (+):** Cu − 2e⁻ → Cu²⁺ (copper dissolves) The anode gradually dissolves, supplying Cu²⁺ ions, while pure copper deposits on the cathode. --- **Question 3:** What is the purpose of electroplating? Give two real-life applications. **Answer:** Electroplating deposits a thin layer of a precious or corrosion-resistant metal onto a base metal object. This improves appearance and prevents rust. **Applications:** (1) Chrome plating on car bumpers to prevent corrosion and enhance shine. (2) Silver plating on tableware and ornaments for aesthetic value and durability. --- **Question 4:** During electrolysis of dilute hydrochloric acid, hydrogen gas is released at the cathode but not chlorine gas. Explain why. **Answer:** In dilute HCl solution, both H⁺ and Cl⁻ ions are present, along with H₂O. At the cathode, H⁺ ions and water molecules compete for electrons. Since H⁺ ions have a lower reduction potential than Cl⁻ ions, H⁺ is reduced preferentially: 2H⁺ + 2e⁻ → H₂. Chlorine is released at the anode instead, where Cl⁻ ions are oxidized more easily than water. --- **Question 5:** Explain the difference between electrolysis and electroplating with reference to electrodes and electrodes used. **Answer:** | Aspect | Electrolysis | Electroplating | |--------|--------------|----------------| | **Purpose** | Break compound into elements or new compounds | Coat one metal with another | | **Anode** | Inert (platinum, graphite) or active | Active (coating metal) | | **Cathode** | Object to be broken down | Object to be coated | | **Electrolyte** | Solution of the compound | Salt solution of coating metal | | **Products** | Gases or new compounds | Thin metal coating on cathode |

**Question 1:** A strip of copper is placed in 100 mL of copper sulphate solution. When electric current is passed through the solution using two copper strips as electrodes, the mass of the copper cathode increases by 3.2 g. Calculate the mass loss at the anode. **Solution:** At the cathode: Cu²⁺ + 2e⁻ → Cu (mass increases) At the anode: Cu − 2e⁻ → Cu²⁺ (mass decreases) Molar mass of Cu = 64 g/mol Moles of Cu deposited at cathode = 3.2 ÷ 64 = 0.05 mol From the equation, for every 1 mole of Cu deposited at cathode, 1 mole of Cu is oxidized at anode. Moles of Cu oxidized at anode = 0.05 mol Mass loss at anode = 0.05 × 64 = **3.2 g** The anode loses the same mass that the cathode gains because identical reactions (in opposite directions) occur. --- **Question 2:** Explain why a solution of dilute sulphuric acid conducts electricity but pure sulphuric acid does not. **Solution:** Pure sulphuric acid (H₂SO₄) is a covalent compound. Although it is a strong acid, the pure liquid contains mostly neutral molecules with very few free ions, so it cannot conduct electricity. When dilute sulphuric acid is prepared (H₂SO₄ added to water), ionization occurs: H₂SO₄ → 2H⁺ + SO₄²⁻ (complete dissociation) The high concentration of free ions (H⁺ and SO₄²⁻) allows electric charge carriers to move freely through the solution. These ions migrate toward opposite electrodes, establishing an electric current. Thus, dilute sulphuric acid is an excellent electrolyte and conducts electricity well. --- **Question 3:** When an inert electrode (graphite) is used in the electrolysis of copper sulphate solution, write the equations for reactions at both electrodes and identify the products. **Solution:** **At the cathode (−, reduction):** Cu²⁺ + 2e⁻ → Cu **Product:** Reddish copper metal deposits on the cathode. **At the anode (+, oxidation):** 2H₂O − 4e⁻ → O₂ + 4H⁺ (Water is oxidized because graphite does not react; Cu²⁺ cannot be further oxidized) **Product:** Oxygen gas bubbles are released. **Overall reaction in solution:** Cu²⁺ + 2H₂O → Cu + O₂ + 4H⁺ (in presence of electric current) Note: The solution becomes more acidic (pH decreases) due to H⁺ ion formation. The mass of the cathode increases while the anode remains unchanged (being inert). --- **Question 4:** In electroplating, why is the object to be plated connected to the cathode and not the anode? Explain with an example. **Solution:** The object is connected to the cathode (negative electrode) because the desired metal must be *deposited* (reduced) onto it. At the cathode, cations of the coating metal are reduced: M⁺ⁿ + ne⁻ → M (metal deposits) If the object were connected to the anode, it would be oxidized and would dissolve instead of receiving a coating. **Example: Silver-plating a copper spoon** - **Cathode:** Copper spoon (object to be plated) - **Anode:** Silver rod or silver plate - **Electrolyte:** Silver nitrate (AgNO₃) solution At the cathode (spoon): Ag⁺ + e⁻ → Ag (silvery coating deposits on spoon) At the anode (silver rod): Ag − e⁻ → Ag⁺ (silver rod dissolves, replenishing Ag⁺ ions) This ensures the spoon receives a uniform, adherent silver coating without being damaged.

**Question 1:** Describe the process of electroplating. What are its industrial applications? Explain with a diagram description how electroplating of silver is performed. **Full Solution:** **Electroplating Definition:** Electroplating is an electrolytic process in which a thin layer of a more precious or corrosion-resistant metal is deposited on the surface of another metal (base metal) by passing electric current through an electrolytic cell. **Principle:** The process relies on the principle of reduction at the cathode, where metal cations from the electrolyte are reduced to metal atoms and deposit on the object. **Apparatus for Silver Electroplating:** - **Anode:** Pure silver plate or rod - **Cathode:** Article (e.g., spoon, jewellery) to be silver-plated - **Electrolyte:** Silver nitrate (AgNO₃) solution (0.5–5 M) or potassium silver cyanide solution - **Power source:** Direct current battery or DC power supply (3–6 V) - **Circuit components:** Ammeter to monitor current, connecting wires **Electrochemical Reactions:** *At the cathode (article to be plated):* Ag⁺(aq) + e⁻ → Ag(s) (Silver ions from solution are reduced to metallic silver and deposit as a shiny coating) *At the anode (silver plate):* Ag(s) − e⁻ → Ag⁺(aq) (Silver atoms from the anode are oxidized, releasing Ag⁺ ions into solution) **Process Steps:** 1. Clean the article thoroughly to remove grease and oxides. 2. Suspend the cleaned article as the cathode. 3. Suspend a pure silver plate as the anode. 4. Immerse both in silver nitrate solution. 5. Connect to DC power (anode to positive, cathode to negative). 6. Pass current at controlled intensity (typically 1–5 amperes, depending on surface area). 7. Silver deposits evenly on the cathode over 30 minutes to several hours. 8. Remove the article and rinse with distilled water. 9. Dry gently to reveal the shiny silver-plated surface. **Industrial Applications:** 1. **Jewellery:** Silver-plating on brass and copper ornaments for shine and protection. 2. **Tableware:** Cutlery, spoons, forks coated with silver for corrosion resistance and aesthetic appeal. 3. **Electronics:** Copper circuit boards plated with tin or gold to improve conductivity and prevent oxidation. 4. **Automotive:** Chrome plating on bumpers, trim, and decorative parts to prevent rust and enhance appearance. 5. **Coins and medals:** Precious metal plating for value and durability. 6. **Medical instruments:** Nickel or chromium plating for sterilization resistance and durability. **Advantages:** - Creates uniform, thin, and adherent coatings. - Protects base metal from corrosion. - Improves aesthetic value. - Cost-effective compared to solid precious metal items. - Controllable coating thickness. **Factors Affecting Electroplating Quality:** - Current density (too high causes rough, brittle deposits; too low is slow). - Temperature of electrolyte (25–40°C optimal). - Duration of plating. - Concentration of metal salt in electrolyte. - Purity of anode metal. --- **Question 2:** Explain the electrolysis of water using inert electrodes. Write equations for reactions at both electrodes, identify gases produced, and state their ratio. **Full Solution:** **Apparatus:** - Two inert electrodes (platinum or graphite rods). - Distilled water with a few drops of dilute sulphuric acid (H₂SO₄) or sodium sulphate (Na₂SO₄) to increase conductivity. - DC power source (6–12 V). - Graduated tubes or gas jars to collect gases. **Electrolysis Process:** When electric current is passed through water: *At the cathode (negative electrode):* 2H⁺ + 2e⁻ → H₂(g) (Hydrogen ions from water are reduced to hydrogen gas) Alternatively, in neutral or slightly basic conditions: 2H₂O + 2e⁻ → H₂(g) + 2OH⁻ *At the anode (positive electrode):* 4OH⁻ − 4e⁻ → O₂(g) + 2H₂O (Hydroxide ions are oxidized to oxygen gas) Alternatively, in acidic conditions: 2H₂O − 4e⁻ → O₂(g) + 4H⁺ **Overall reaction:** 2H₂O(l) → 2H₂(g) + O₂(g) (via electrolysis) **Observations:** 1. Two gases are produced: hydrogen (at cathode) and oxygen (at anode). 2. Hydrogen gas burns with a pop sound (test with burning splint). 3. Oxygen supports combustion (glowing splint reignites in oxygen). 4. Oxygen is produced in smaller volume; hydrogen in larger volume. 5. The electrolyte (acidified water) remains unchanged in composition. **Volume Ratio of Gases:** V(H₂) : V(O₂) = 2 : 1 For example, if 200 mL of hydrogen is collected at the cathode, 100 mL of oxygen is collected at the anode. **Molar Ratio:** n(H₂) : n(O₂) = 2 : 1 Mass produced in 5-minute electrolysis: - Hydrogen gas: typically 1–2 mL (mass ≈ 0.002–0.004 g) - Oxygen gas: typically 0.5–1 mL (mass ≈ 0.008–0.016 g, denser than H₂) **Why Add Acid or Salt?** Pure distilled water is a very poor conductor (very low ionic concentration). Adding a few drops of dilute H₂SO₄ or a pinch of Na₂SO₄ increases the concentration of free ions (H⁺, SO₄²⁻ or Na⁺, SO₄²⁻), making the solution sufficiently conductive to allow measurable current flow without interfering with the main reaction (water decomposition). **Practical Significance:** - Industrial hydrogen production via water electrolysis (green hydrogen). - Oxygen obtained is pure and can be used in medical and industrial applications. - Demonstrates the law of conservation of mass and energy. --- **Question 3:** Compare electrolysis and electroplating in terms of purpose, electrodes used, electrolyte composition, and products formed. Present your answer in a table format with explanations. **Full Solution:** | **Parameter** | **Electrolysis** | **Electroplating** | |---|---|---| | **Main Purpose** | Decompose chemical compounds into simpler substances (elements or new compounds); industrial production of chemicals | Coat an object with a thin layer of another metal to improve appearance, corrosion resistance, or durability | | **Anode** | Inert (graphite, platinum) or active (same metal as solution) | Active (the coating metal; must be pure and connected to positive terminal) | | **Cathode** | The substance to be decomposed or an inert electrode | The object to be coated (connected to negative terminal) | | **Electrolyte** | Solution or molten compound to be decomposed; e.g., CuSO₄ solution, molten NaCl, dilute H₂SO₄ | Salt solution of the coating metal; e.g., AgNO₃ for silver plating, CuSO₄ for copper plating, or [Ag(CN)₂]⁻ for silver cyanide plating | | **Reactions at Cathode** | Reduction of cations (or water); e.g., Cu²⁺ + 2e⁻ → Cu | Reduction of coating metal cations; e.g., Ag⁺ + e⁻ → Ag | | **Reactions at Anode** | Oxidation of anions (or water); e.g., 2Cl⁻ − 2e⁻ → Cl₂ or 2H₂O − 4e⁻ → O₂ + 4H⁺ | Oxidation of coating metal; e.g., Ag − e⁻ → Ag⁺ (anode dissolves, replenishing ions) | | **Products Formed** | New substances (often gases or metals); e.g., H₂, O₂, Cl₂, Cu metal, etc. | Thin, uniform metal layer on cathode; anode may dissolve (if active) | | **Duration** | Variable; depends on amount to decompose and current intensity | Controlled; typically 30 minutes to several hours | | **Current Efficiency** | Often <100% (side reactions may occur) | High (>95%); fewer competing reactions | | **Industrial Applications** | Extraction of metals (Al from Al₂O₃), production of chlorine and caustic soda, refining of copper | Decorative coatings (silver, gold jewellery), corrosion protection (chrome bumpers), electronics (tin coating on PCBs) | **Detailed Explanations:** **1. Anode Material:** - In **electrolysis**, if an active anode is used (e.g., copper in CuSO₄ electrolysis), it dissolves: Cu − 2e⁻ → Cu²⁺. If inert (graphite in water electrolysis), it does not react; instead, water or anions oxidize. - In **electroplating**, the anode *must* be active and pure to ensure a steady supply of coating metal ions into the electrolyte. For example, in silver plating, a pure silver anode dissolves at the same rate that silver deposits on the cathode, maintaining Ag⁺ concentration. **2. Electrolyte Role:** - In **electrolysis**, the electrolyte contains the ions to be decomposed. Changing the electrolyte changes the products. For example, electrolysis of NaCl solution gives H₂ and O₂ (water ions are preferentially discharged), whereas electrolysis of molten NaCl gives Na metal and Cl₂ gas. - In **electroplating**, the electrolyte must contain a soluble salt of the coating metal at appropriate concentration. This ensures sufficient metal cations are available for reduction at the cathode. **3. Control and Quality:** - **Electrolysis** is often a decomposition process with less stringent control (though current, voltage, and temperature affect yields). - **Electroplating** requires precise control of current density (typically 1–5 A/dm²), temperature, and electrolyte composition to achieve a uniform, smooth, and adherent coating. Too high current causes rough, porous deposits; too low current is inefficient. **4. Economic Significance:** - **Electrolysis** is essential for bulk chemical and metal production (e.g., 99% of world's aluminium is produced via electrolysis of Al₂O₃). - **Electroplating** adds value to consumer goods and protects expensive items from corrosion, extending their lifespan. Explore more patterns and practice tests with cbsetutor.ai's interactive AI tutor, which drills these question types daily with real-time feedback.

**Case Study: Silver Jewellery Manufacturing** A jewellery manufacturer receives 500 brass brooches (copper–zinc alloy) that must be silver-plated for a premium collection. The electroplating tank contains 2 L of AgNO₃ solution (concentration: 0.5 M). Each brooch has a total surface area of 8 cm². A pure silver anode (area 50 cm²) is used, and a current of 2 A is applied. **Given Data:** - Molar mass of Ag = 108 g/mol - Faraday's constant (F) = 96,500 C/mol - Density of Ag = 10.5 g/cm³ - Required coating thickness = 0.02 mm per brooch **Questions:** **(a)** Calculate the mass of silver deposited on one brooch and the total mass for 500 brooches. **(b)** Estimate the time (in hours) required to plate all 500 brooches. **(c)** Will the AgNO₃ solution be sufficient to plate all 500 brooches? Calculate the moles of Ag⁺ available and moles required. **(d)** Why is the pure silver anode essential? What would happen if a copper anode were used instead? --- **Solution Steps:** **(a) Mass of Silver on Brooches:** **Volume of silver coating per brooch:** Volume = Surface area × Thickness V = 8 cm² × 0.02 mm = 8 cm² × 0.002 cm = 0.016 cm³ **Mass of silver per brooch:** m = Density × Volume m = 10.5 g/cm³ × 0.016 cm³ = 0.168 g ≈ 0.17 g **Total mass for 500 brooches:** Total mass = 0.17 g × 500 = **85 g** --- **(b) Time Required for Electroplating:** **Moles of Ag to be deposited:** n(Ag) = Total mass ÷ Molar mass n(Ag) = 85 g ÷ 108 g/mol ≈ 0.787 mol **Reduction reaction at cathode:** Ag⁺ + e⁻ → Ag Each mole of Ag requires 1 mole of electrons. **Total charge (coulombs) required:** Q = n × F = 0.787 mol × 96,500 C/mol = 75,955.5 C **Time calculation using Q = I × t:** t = Q ÷ I t = 75,955.5 C ÷ 2 A = 37,977.75 seconds **Convert to hours:** t = 37,977.75 s ÷ 3,600 s/h ≈ **10.5 hours** (or approximately 10 hours 30 minutes) --- **(c) Sufficiency of AgNO₃ Solution:** **Moles of Ag⁺ available in the tank:** n(Ag⁺) = Concentration × Volume n(Ag⁺) = 0.5 M × 2 L = **1 mol** **Moles of Ag⁺ required for plating:** From part (a): n(Ag⁺) required = 0.787 mol **Comparison:** Available = 1 mol Required = 0.787 mol **Yes, the solution is sufficient** because 1 mol > 0.787 mol. There is approximately 0.213 mol (or 23 g) of excess Ag⁺ remaining. However, the solution concentration will drop as silver is deposited, which may slow the process. The pure silver anode also continuously dissolves, releasing Ag⁺: Ag − e⁻ → Ag⁺ This replenishes the electrolyte, ensuring adequate ion concentration throughout plating. --- **(d) Role of Pure Silver Anode & Consequences of Copper Anode:** **Why the pure silver anode is essential:** 1. **Ion replenishment:** As Ag⁺ ions are reduced at the cathode, the silver anode dissolves to replenish Ag⁺ in solution, maintaining constant electrolyte concentration and current efficiency. 2. **Ensures uniform coating:** Consistent Ag⁺ supply produces smooth, even deposition across all brooches. 3. **Prevents electrode polarization:** Without anode dissolution, Ag⁺ concentration would drop sharply, requiring higher voltage to sustain current, causing side reactions and rough deposits. **If a copper anode were used:** 1. **Wrong reaction at anode:** Instead of dissolving and releasing Ag⁺, copper would oxidize: Cu − 2e⁻ → Cu²⁺ 2. **Ag⁺ depletion:** The electrolyte would rapidly lose Ag⁺ without replenishment, reducing plating efficiency and increasing voltage drop. 3. **Contamination:** Cu²⁺ ions would mix with the electrolyte, potentially depositing copper onto brooches instead of (or mixed with) silver, ruining the quality. 4. **Uneven coating:** As Ag⁺ depletes, cathode current density decreases unevenly, producing thin, patchy silver layers. 5. **Increased time and cost:** Much higher voltage and longer plating time would be needed, making the process uneconomical. **Conclusion:** The pure silver anode is indispensable for efficient, uniform, high-quality electroplating. Material selection for electrodes is as critical as circuit design in electrochemistry. --- **Follow-up Thinking Questions:** - How would increasing current to 4 A affect plating time? (Answer: Time halved to ~5.25 hours) - If brooches are arranged in series vs. parallel, how would current distribution differ? (Series reduces current per brooch; parallel keeps current constant) - What would happen if the solution temperature rose to 50°C? (Faster ion diffusion, faster plating, but risk of burnt or rough deposits)

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