Chapter 1: Chemical Reactions and Equations - Practice worksheet with complete answer key
According to your class and board
- Board: CBSE
- Class/Year: 10
- Subject: Chemistry
- Book + Publisher: Science (NCERT)
- Chapter: Chemical Reactions and Equations
- Difficulty: Standard
- Goal: Practice
A1. Define a chemical reaction.
A2. What is meant by a 'balanced chemical equation'?
A3. Name the substance that gets oxidised in the reaction: CuO + H₂ → Cu + H₂O.
A4. What is a precipitate?
A5. Give one example of a decomposition reaction.
B1. Balance the equation: Fe + H₂O → Fe₃O₄ + H₂
B2. Distinguish between a combination reaction and a decomposition reaction with one example each.
B3. What is a redox reaction? Identify the oxidising and reducing agents in: ZnO + C → Zn + CO.
B4. Why is respiration considered an exothermic reaction?
B5. Write the balanced equation for the reaction between sodium and water.
B6. What type of reaction is represented by: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)?
B7. Why do we store silver chloride in dark-coloured bottles?
B8. What happens when ferrous sulphate crystals are heated? Write the equation.
B9. Define 'rancidity'. How can it be prevented?
B10. What is the role of sunlight in the decomposition of silver bromide?
C1. Iron articles are painted to prevent rusting. Explain the chemistry behind this preventive measure.
C2. A white compound (X) on heating gives a yellow residue and a gas that turns lime water milky. Identify X and write the reaction.
C3. Why are chips packets flushed with nitrogen gas? Relate your answer to the type of chemical reaction involved.
C4. When barium chloride solution is mixed with sodium sulphate solution, a white precipitate is formed. Identify the precipitate and write the balanced equation.
C5. A student adds dilute HCl to a test tube containing zinc granules. What will he observe? Write the equation.
C6. Why does the colour of copper sulphate solution change when an iron nail is dipped into it? Identify the type of reaction.
C7. Food becomes stale after a few days. Name the chemical process and suggest two methods to slow it down.
C8. Decomposition of vegetable matter into compost is an example of which type of reaction? Is it exothermic or endothermic?
D1. What is corrosion? Explain the conditions necessary for rusting of iron. Suggest three methods to prevent it. (3 marks)
D2. Classify the following reactions and balance them: (a) BaCl₂ + Na₂SO₄ → BaSO₄ + NaCl (b) 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂. (3 marks)
D3. In the reaction: MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂, identify the substance oxidised, the substance reduced, the oxidising agent, and the reducing agent. (3 marks)
D4. What are exothermic and endothermic reactions? Give two examples of each from daily life. (3 marks)
D5. Write the equation for the reaction of quicklime with water. What type of reaction is it? Is it exothermic or endothermic? (2 marks)
E1. A student writes: Na + O₂ → Na₂O as a balanced equation. Is this correct?
E2. "All decomposition reactions are endothermic." Is this statement true? Justify.
A1. A chemical reaction is a process in which one or more substances (reactants) are transformed into new substances (products) with different properties.
A2. A balanced chemical equation has an equal number of atoms of each element on both the reactant and product sides.
A3. H₂ is oxidised (gains oxygen → becomes H₂O). CuO is reduced.
A4. A precipitate is an insoluble solid formed during a chemical reaction in solution.
A5. 2H₂O → 2H₂ + O₂ (electrolysis of water).
B1. 3Fe + 4H₂O → Fe₃O₄ + 4H₂.
B2. Combination: Two or more substances combine to form a single product (e.g., 2Mg + O₂ → 2MgO). Decomposition: A single compound breaks into two or more simpler substances (e.g., 2HgO → 2Hg + O₂).
B3. A redox reaction involves both oxidation and reduction simultaneously. In ZnO + C → Zn + CO: C is the reducing agent (gets oxidised to CO), ZnO is the oxidising agent (Zn is reduced).
B4. Respiration releases energy: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy. Since energy is released, it is exothermic.
B5. 2Na + 2H₂O → 2NaOH + H₂↑.
B6. Double displacement reaction (also a precipitation reaction since AgCl is an insoluble precipitate).
B7. Silver chloride decomposes in sunlight: 2AgCl → 2Ag + Cl₂. Dark bottles prevent this photolytic decomposition.
B8. Green ferrous sulphate crystals turn white (lose water), then decompose: 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃. The residue is reddish-brown.
B9. Rancidity is the oxidation of fats and oils in food, causing bad smell and taste. Prevention: store in airtight containers, add antioxidants, refrigerate, flush with nitrogen.
B10. Sunlight provides the energy needed for decomposition: 2AgBr → 2Ag + Br₂. This is a photochemical decomposition reaction.
C1. Paint creates a barrier between iron and moisture/oxygen, preventing the electrochemical reaction (4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O) that causes rusting. Without contact with both water and air, rusting cannot occur.
C2. X is zinc carbonate (ZnCO₃). ZnCO₃ → ZnO (yellow) + CO₂↑ (turns lime water milky). It is a thermal decomposition reaction.
C3. Nitrogen is an inert gas. Flushing packets with N₂ prevents oxidation (a slow chemical reaction) of fats in chips, which would cause rancidity. It displaces oxygen and moisture.
C4. Precipitate: BaSO₄ (white). BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s)↓ + 2NaCl(aq). This is a double displacement reaction.
C5. Observation: Zinc dissolves, colourless gas (H₂) evolves with brisk effervescence. Zn + 2HCl → ZnCl₂ + H₂↑. Displacement reaction.
C6. Iron displaces copper from the solution: Fe + CuSO₄ → FeSO₄ + Cu. The blue solution turns green (FeSO₄) and reddish-brown copper deposits on the nail. It is a displacement reaction.
C7. Oxidation / decomposition. Methods: refrigeration, adding preservatives, vacuum packing, using antioxidants.
C8. Exothermic decomposition reaction (heat is released during composting, which is why compost heaps feel warm).
D1. Corrosion: Surface degradation of metals by reaction with air, moisture, or chemicals. For rusting: iron needs both water and oxygen. 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O. Prevention: (i) Painting/oiling (ii) Galvanisation (zinc coating) (iii) Electroplating or alloying (stainless steel).
D2. (a) BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl — Double displacement reaction. (b) Already balanced: 2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂ — Thermal decomposition reaction.
D3. HCl is oxidised (Cl⁻ → Cl₂, loses electrons). MnO₂ is reduced (Mn⁴⁺ → Mn²⁺, gains electrons). Oxidising agent: MnO₂. Reducing agent: HCl.
D4. Exothermic (release heat): (i) Burning of fuels (ii) Respiration. Endothermic (absorb heat): (i) Photosynthesis (ii) Dissolving ammonium chloride in water.
D5. CaO + H₂O → Ca(OH)₂ + Heat. Combination reaction. It is exothermic (large amount of heat is produced — the mixture gets very hot).
E1. No. The balanced equation is: 4Na + O₂ → 2Na₂O. The student didn't balance the sodium atoms (4 on the right, only 1 on the left in the original). — Why students get this wrong: They forget to count atoms inside compound formulas and assume a 1:1 ratio.
E2. No, this is false. While many decomposition reactions require energy (endothermic), some are exothermic. For example, decomposition of vegetable matter (composting) releases heat. Decomposition of hydrogen peroxide (2H₂O₂ → 2H₂O + O₂) is also exothermic. — Why students get this wrong: They associate "breaking bonds" with "requiring energy" but forget that the products may release more energy than was needed to break the reactant bonds.