Practice Questions for Science Class 10th "Chemical Reactions and Equations"

Multiple Choice Questions (MCQs):

1. B) Burning of paper - Burning involves a chemical change where new substances are formed.
2. B) Reactant - Reactants are the substances that undergo change in a chemical reaction.
3. A) Must be equal on both sides - A balanced equation ensures the conservation of mass by having the same number of atoms on both sides.
4. A) CaO + H₂O → Ca(OH)₂ - This is a combination reaction where two substances combine to form a single compound.
5. A) Oxidation - Adding oxygen to a substance is one definition of oxidation.
6. B) Endothermic - Heat absorbed reactions are endothermic.
7. A) Aqueous solution - (aq) denotes that the substance is in an aqueous (water) solution.
8. D) Both B and C - Photosynthesis is both endothermic (absorbs light energy) and a combination reaction (CO₂ + H₂O → C₆H₁₂O₆ + O₂).
9. A) CaO + CO₂ - When heated, calcium carbonate decomposes to calcium oxide and carbon dioxide.
10. C) Burning of natural gas - Exothermic reactions release heat, and combustion of natural gas is a prime example.
11. A) The exchange of ions between two compounds - In double displacement, ions from two compounds switch places.
12. A) A displacement reaction - Sodium displaces hydrogen in water: 2Na + 2H₂O → 2NaOH + H₂.
13. A) Oxidized - Magnesium gains oxygen, hence it is oxidized.
14. D) Melting of a solid - Melting is a physical change, not a sign of a chemical reaction.
15. A) A redox reaction - Rusting involves oxidation of iron and reduction of oxygen.
16. C) It provides an alternate reaction pathway with a lower activation energy - Catalysts speed up reactions without being consumed.
17. D) Double displacement - Ions from AgNO₃ and NaCl exchange to form AgCl and NaNO₃.
18. B) Zn + 2HCl → ZnCl₂ + H₂ - This is the correct balanced equation for the reaction.
19. C) Both oxidation and reduction - Rusting involves iron oxidation and oxygen reduction.
20. B) Exothermic - If the product energy is less than the reactant energy, heat is released.

Short Answer Questions:

21. Chemical Equation: A symbolic representation of a chemical reaction where reactants are shown on the left and products on the right.
22. Balanced Chemical Equation: An equation where the number of atoms for each element is the same on both sides, reflecting the conservation of mass.
23. Exothermic vs. Endothermic:

• Exothermic: Releases heat (e.g., burning of fuel).
• Endothermic: Absorbs heat (e.g., photosynthesis).

24. Oxidation: Involves the loss of electrons or the gain of oxygen by a substance.
25. Types of Chemical Reactions:

• Combination: A + B → AB (e.g., 2H₂ + O₂ → 2H₂O)
• Decomposition: AB → A + B (e.g., 2H₂O → 2H₂ + O₂)
• Displacement: A + BC → B + AC (e.g., Zn + CuSO₄ → ZnSO₄ + Cu)
• Double Displacement: AB + CD → AD + CB (e.g., AgNO₃ + NaCl → AgCl + NaNO₃)
• Redox: Involves electron transfer (e.g., rusting of iron).

26. Recognizing Chemical Change: Change in color, temperature, odor, formation of a gas or precipitate, or release/absorption of energy.
27. Precipitation Reaction: When two solutions react to form an insoluble solid. Example: AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq).
28. Magnesium Burning in Air: 2Mg + O₂ → 2MgO
29. Corrosion: The natural process of metal degradation through oxidation, e.g., rusting of iron (4Fe + 3O₂ → 2Fe₂O₃).
30. Catalyst Function: Provides an alternative pathway with lower activation energy, speeding up the reaction without being consumed.
31. Balancing Equations: Ensures the law of conservation of mass is upheld, showing that no atoms are created or destroyed.
32. Law of Conservation of Mass: The total mass of reactants equals the total mass of products in a closed system.
33. Respiration as a Reaction: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy (ATP), an exothermic reaction.
34. Acids/Bases and Reaction Rate: They can change the pH, affecting the rate by altering the concentration of reactive ions or by catalyzing reactions.
35. Displacement Reaction: One element displaces another from a compound. Example: Zn + CuSO₄ → ZnSO₄ + Cu.
36. Physical vs. Chemical Change:

• Physical: No new substance formed, reversible (e.g., melting ice).
• Chemical: New substances formed, often irreversible (e.g., rusting).

37. Temperature and Reaction Rate: Higher temperatures increase reaction rates by providing more energy to overcome activation energy.
38. State Symbols in Equations: Indicate the physical state of reactants and products (s - solid, l - liquid, g - gas, aq - aqueous), crucial for understanding reaction conditions.
39. Calcium Oxide and Water: CaO + H₂O → Ca(OH)₂, products are calcium hydroxide and heat (an exothermic reaction).
40. Metals Not Reacting with Acids: Less reactive metals like copper or silver do not displace hydrogen from acids, as they are below hydrogen in the reactivity series.

Long Answer Questions:

41. Types of Chemical Reactions:

• Combination: Two or more substances combine.
• Decomposition: A single compound breaks down.
• Displacement: One element displaces another in a compound.
• Double Displacement: Ions from two compounds switch places.
• Redox: Involves oxidation-reduction where electrons are transferred.

42. Balancing Equations:

• Example: H₂ + O₂ → H₂O
• Steps:
  1. Count atoms: H = 2, O = 2 on reactant side; H = 2, O = 1 on product side.
  2. Balance O first: H₂ + O₂ → 2H₂O
  3. Balance H: 2H₂ + O₂ → 2H₂O

43. Oxidation and Reduction:

• Oxidation: Loss of electrons or gain of oxygen.
• Reduction: Gain of electrons or loss of oxygen.
• Redox Reaction: Both processes occur simultaneously, e.g., in rusting.

44. Factors Affecting Reaction Rate:

• Temperature: Higher temperature, faster rate (e.g., cooking).
• Concentration: More concentrated reactants speed up reactions (e.g., acid-base neutralization).
• Catalyst: Speeds up reaction by lowering activation energy (e.g., enzymes in biological systems).

45. Combination vs. Decomposition:

• Combination: Two or more substances combine (e.g., synthesis of water).
• Decomposition: One compound breaks down into simpler substances (e.g., electrolysis of water).

46. Chemical Reactions in Daily Life:

• Cooking: Maillard reaction for flavor.
• Photosynthesis: Plants converting CO₂ and H₂O to glucose.
• Respiration: Human body converting glucose to energy.

47. Reversible Reactions:

• A reaction where products can revert to reactants under certain conditions.
• Example: Haber process for ammonia (N₂ + 3H₂ ⇄ 2NH₃).

48. Preventing Rust:

• Chemical: Applying inhibitors or sacrificial protection.
• Physical: Painting, galvanization, or using rust-resistant alloys.

49. Activity Series:

• A list of metals ordered by their reactivity, used to predict displacement reactions. If one metal is above another, it can displace it from its compounds.

50. Stoichiometry and Equations:

• Chemical equations show reactant-product ratios, crucial for calculating quantities in reactions, determining limiting reactants, and predicting yields.

51. Pressure and Reaction Rate:

• Increases rate in gas reactions by increasing reactant concentration, e.g., in the synthesis of ammonia.

52. Identifying Unknown Substances:

• Reactions like precipitation or color change can identify ions or compounds, e.g., using silver nitrate to detect chloride ions.

53. Environmental Impacts:

• Combustion: Produces CO₂, contributing to greenhouse effect.
• Corrosion: Leads to material loss and environmental contamination if not managed.

54. Limiting Reactants:

• The reactant that is completely consumed first, determining the amount of product. Example: in N₂ + 3H₂ → 2NH₃, if there's less N₂, it limits NH₃ production.

55. Exothermic/Endothermic in Industry:

• Exothermic: Energy efficiency, less energy input needed.
• Endothermic: Requires energy but can be used

55. Exothermic/Endothermic in Industry (Continued):

• Exothermic: Helps in processes like welding or the manufacture of cement where heat is beneficial.
• Endothermic: Useful in processes requiring cooling or where products need to be at a higher energy state, like in the production of nitrogenous fertilizers.

56. Enzymes as Biological Catalysts:

• Enzymes are proteins that act as catalysts in biological systems, speeding up reactions like digestion:
  • Amylase: Converts starch to sugars in the mouth and small intestine.
  • Pepsin: Breaks down proteins in the stomach.
• They lower the activation energy, allowing reactions to occur at body temperature.

57. Chemical Reactions in Environmental Management:

• Waste Treatment: Bioremediation uses microbial reactions to degrade pollutants.
• Water Purification: Chemical reactions like chlorination or ozonation for disinfecting water.
• Air Pollution Control: Catalytic converters in vehicles use redox reactions to convert harmful gases into less harmful substances.

58. Energy Changes in Chemical Reactions:

• Activation Energy: The initial energy needed to start a reaction, often reduced by catalysts.
• Enthalpy Change (ΔH):
  • Positive ΔH for endothermic reactions (energy absorbed).
  • Negative ΔH for exothermic reactions (energy released).
• Understanding these helps in managing reaction conditions for industrial processes.

59. Significance of Reaction Rate in Manufacturing:

• Affects production speed, quality control, and safety.
• Example: In pharmaceutical production, reaction rates must be controlled to ensure drug purity and yield.

60. Chemical Reactions in Materials Science:

• Polymers, ceramics, and alloys are developed through specific chemical reactions:
  • Polymerization: Creating long-chain molecules for plastics.
  • Alloying: Mixing metals to achieve desired properties.
• Understanding reactions allows for the design of materials with specific characteristics for applications in electronics, construction, etc.

Application-Based Questions (Continued):

61. Limiting Reactant in Water Formation:

• 2H₂ + O₂ → 2H₂O
• With 4 moles of H₂ and 2 moles of O₂, neither is limiting; they match the stoichiometric ratio for complete reaction, producing 4 moles of H₂O.

62. Experiment for Mass Conservation:

• Conduct a simple reaction in a sealed container, like the reaction between baking soda (NaHCO₃) and vinegar (acetic acid, HC₂H₃O₂), measure mass before and after; it should remain constant.

63. Demonstrating Endothermic Reaction:

• Use the dissolution of ammonium nitrate in water. Place a thermometer in the solution, noting the temperature drop as the reaction absorbs heat from the surroundings.

64. Barium Chloride and Sodium Sulfate Reaction:

• BaCl₂ (aq) + Na₂SO₄ (aq) → BaSO₄ (s) + 2NaCl (aq)
• The precipitate is Barium Sulfate (BaSO₄).

65. Burning Methane:

• CH₄ + 2O₂ → CO₂ + 2H₂O
• Types of reaction: Combustion (burning in oxygen) and oxidation.

66. Mass of CO₂ from CaCO₃ Decomposition:

• CaCO₃ → CaO + CO₂
• Moles of CaCO₃ = 10g / 100g/mol = 0.1 mol
• Mass of CO₂ = 0.1 mol × 44 g/mol = 4.4 grams

Critical Thinking Questions:

67. Chemical Reactions and Pollution:

• Combustion of fossil fuels releases CO₂, contributing to the greenhouse effect; corrosion can lead to metal ions polluting water bodies.

68. Limiting Reactants and Reaction Yield:

• Example: N₂ + 3H₂ → 2NH₃. If 3 moles of N₂ and 6 moles of H₂ are used, N₂ limits the reaction, allowing only 2 moles of NH₃ to form, not 4.

69. Exothermic vs. Endothermic in Industry:

• Exothermic: Allows for energy savings, as seen in welding or cement production.
• Endothermic: Useful in processes where heat management is critical, like in the endothermic step of ammonia synthesis where heat must be added.

70. Enzymes in Digestion:

• Enzymes like lipase, protease, and amylase catalyze the breakdown of fats, proteins, and carbohydrates, respectively, enhancing digestion efficiency at body temperature.

71. Chemical Reactions in Environmental Management:

• Using reactions to convert pollutants into harmless compounds or to precipitate harmful ions out of waste streams.

72. Energy Changes in Reactions:

• Activation energy is the energy barrier to initiate a reaction; catalysts lower this barrier. Enthalpy changes indicate whether the reaction releases or absorbs heat.

73. Rate of Reaction in Manufacturing:

• Critical for controlling production processes, ensuring product consistency, and optimizing resource use.

74. Chemical Reactions and Materials Science:

• Reactions are used to synthesize materials with specific properties, like the vulcanization of rubber or the synthesis of semiconductors.

75. Pressure's Effect on Reaction Rate:

• For gas-phase reactions, pressure increases rate by increasing concentration, as seen in high-pressure processes like the production of ammonia.

76. Identifying Unknown Substances:

• By observing unique reaction characteristics like color changes, gas evolution, or precipitate formation, one can deduce the presence of certain ions or compounds.

77. Environmental Impacts of Reactions:

• Combustion contributes to air pollution and climate change, while corrosion leads to material degradation and potential environmental contamination.

78. Limiting Reactants and Yield:

• Understanding limiting reactants is crucial for predicting how much product can be formed, which is vital in industrial settings for cost and efficiency.

79. Exothermic and Endothermic in Industry:

• Exothermic reactions can be harnessed for heat generation, while endothermic ones might require careful energy management.

80. Chemical Reactions and Materials Development:

• From creating new polymers to developing advanced ceramics or alloys, chemical reactions are at the heart of materials innovation, influencing everything from consumer products to aerospace materials.