Class 10 Notes

Class 10 Science Notes Chapter 1 (Chapter 1) – Examplar Problems (English) Book

Philoid

Philoid

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Examplar Problems (English)
Detailed Notes with MCQs of Chapter 1: Chemical Reactions and Equations from your NCERT Exemplar. This chapter is fundamental, not just for your board exams, but also forms the basis for many questions in competitive government exams. Pay close attention to the details.

Chapter 1: Chemical Reactions and Equations - Detailed Notes for Competitive Exams

1. Chemical Reaction:

  • Definition: A process involving the rearrangement of the structure of molecules or ions into new substances, different from the initial ones. It involves the breaking of old chemical bonds and the formation of new chemical bonds.
  • Indicators/Characteristics: How do we know a chemical reaction has occurred? Look for these signs:
    • Change in State: Solid to liquid/gas, liquid to gas, etc. (e.g., Burning of candle wax - solid to liquid and gas).
    • Change in Colour: (e.g., Iron rusting - silvery grey to reddish-brown; Copper sulphate reacting with iron - blue to pale green).
    • Evolution of a Gas: (e.g., Zinc reacting with dilute sulphuric acid produces hydrogen gas; Reaction of sodium carbonate with dilute HCl produces carbon dioxide).
    • Change in Temperature:
      • Exothermic Reaction: Heat is released, temperature increases (e.g., Reaction of quicklime (CaO) with water; Respiration).
      • Endothermic Reaction: Heat is absorbed, temperature decreases (e.g., Decomposition of calcium carbonate; Reaction of barium hydroxide with ammonium chloride).
    • Formation of a Precipitate: An insoluble solid formed during a reaction in a solution (e.g., Reaction between lead nitrate and potassium iodide forms a yellow precipitate of lead iodide (PbI₂)).

2. Chemical Equation:

  • Definition: A symbolic representation of a chemical reaction using chemical formulae and symbols.
  • Word Equation: Represents reactants and products using their names (e.g., Magnesium + Oxygen → Magnesium oxide).
  • Skeletal Chemical Equation: Uses formulae but is unbalanced (e.g., Mg + O₂ → MgO).
  • Balanced Chemical Equation: The number of atoms of each element is equal on both the reactant and product sides. This adheres to the Law of Conservation of Mass (mass can neither be created nor destroyed in a chemical reaction).
    • Balancing Method (Hit and Trial): Adjust stoichiometric coefficients (numbers placed before formulae) to equalize the atom count for each element on both sides. Never change the chemical formula itself.
    • Example: Balancing Fe + H₂O → Fe₃O₄ + H₂
      • Fe + H₂O → Fe₃O₄ + H₂ (Skeletal)
      • Count atoms: Fe(1), H(2), O(1) → Fe(3), H(2), O(4)
      • Balance Fe: 3Fe + H₂O → Fe₃O₄ + H₂
      • Balance O: 3Fe + 4H₂O → Fe₃O₄ + H₂
      • Balance H: 3Fe + 4H₂O → Fe₃O₄ + 4H₂
      • Check: Fe(3), H(8), O(4) → Fe(3), H(8), O(4) (Balanced)
  • Making Equations More Informative:
    • Physical States: (s) for solid, (l) for liquid, (g) for gas, (aq) for aqueous solution.
      • Example: 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g) (Note: Steam is used, hence H₂O(g))
    • Reaction Conditions: Temperature, pressure, catalyst written above or below the arrow (e.g., sunlight, heat (Δ)).
    • Heat Changes: Indicate exothermic (+ Heat or ΔH = -ve) or endothermic (- Heat or ΔH = +ve) nature.

3. Types of Chemical Reactions:

  • a) Combination Reaction:

    • Definition: Two or more reactants combine to form a single product.
    • General Form: A + B → AB
    • Examples:
      • Burning of coal: C(s) + O₂(g) → CO₂(g) + Heat
      • Formation of water: 2H₂(g) + O₂(g) → 2H₂O(l) + Heat
      • Slaking of lime: CaO(s) (Quicklime) + H₂O(l) → Ca(OH)₂(aq) (Slaked lime) + Heat (Highly exothermic)
      • Note: Ca(OH)₂ solution (limewater) reacts with CO₂ to form CaCO₃ (white precipitate), causing milkiness. Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l). Excess CO₂ dissolves the precipitate: CaCO₃(s) + H₂O(l) + CO₂(g) → Ca(HCO₃)₂(aq) (Calcium bicarbonate - soluble).

  • b) Decomposition Reaction:

    • Definition: A single reactant breaks down into two or more simpler products. Usually requires energy (heat, light, electricity). Mostly endothermic.
    • General Form: AB → A + B
    • Types:
      • Thermal Decomposition: Decomposition by heat.
        • FeSO₄·7H₂O(s) (Green vitriol) --Heat--> FeSO₄(s) (White) + 7H₂O(g)
        • 2FeSO₄(s) (Ferrous sulphate) --Strong Heat--> Fe₂O₃(s) (Ferric oxide - brown) + SO₂(g) + SO₃(g) (Sulphur dioxide & trioxide - pungent smell)
        • CaCO₃(s) (Limestone) --Heat (Δ)--> CaO(s) (Quicklime) + CO₂(g) (Used in cement industry)
        • 2Pb(NO₃)₂(s) (Lead nitrate - white) --Heat (Δ)--> 2PbO(s) (Lead oxide - yellow) + 4NO₂(g) (Nitrogen dioxide - brown fumes) + O₂(g) (Oxygen)
      • Electrolytic Decomposition (Electrolysis): Decomposition by passing electric current.
        • 2H₂O(l) --Electric Current--> 2H₂(g) + O₂(g) (Hydrogen collects at cathode, Oxygen at anode; Volume of H₂ is double that of O₂)
        • 2NaCl(molten) --Electric Current--> 2Na(s) + Cl₂(g)
      • Photolytic Decomposition (Photolysis): Decomposition by light.
        • 2AgCl(s) (Silver chloride - white) --Sunlight--> 2Ag(s) (Silver - grey) + Cl₂(g) (Chlorine - yellowish-green gas)
        • 2AgBr(s) (Silver bromide - pale yellow) --Sunlight--> 2Ag(s) + Br₂(g) (Bromine - reddish-brown gas) (Used in black and white photography)

  • c) Displacement Reaction:

    • Definition: A more reactive element displaces a less reactive element from its compound (usually in solution). Reactivity is determined by the Reactivity Series.
    • General Form: A + BC → AC + B (where A is more reactive than B)
    • Reactivity Series (simplified): K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au > Pt (Most reactive to Least reactive)
    • Examples:
      • Fe(s) + CuSO₄(aq) (Blue) → FeSO₄(aq) (Pale Green) + Cu(s) (Reddish-brown deposit) (Fe is more reactive than Cu)
      • Zn(s) + CuSO₄(aq) (Blue) → ZnSO₄(aq) (Colourless) + Cu(s)
      • Pb(s) + CuCl₂(aq) (Greenish-blue) → PbCl₂(aq) (Colourless) + Cu(s)
      • Cu(s) + AgNO₃(aq) (Colourless) → Cu(NO₃)₂(aq) (Blue) + 2Ag(s) (Shiny silver deposit) (Cu is more reactive than Ag)
      • Note: A less reactive metal cannot displace a more reactive metal (e.g., Cu(s) + FeSO₄(aq) → No Reaction).

  • d) Double Displacement Reaction:

    • Definition: A reaction in which there is an exchange of ions between the reactants.
    • General Form: AB + CD → AD + CB
    • Types:
      • Precipitation Reaction: A reaction that produces an insoluble solid (precipitate).
        • Na₂SO₄(aq) (Sodium sulphate) + BaCl₂(aq) (Barium chloride) → BaSO₄(s) (Barium sulphate - white precipitate) + 2NaCl(aq) (Sodium chloride)
        • Pb(NO₃)₂(aq) (Lead nitrate) + 2KI(aq) (Potassium iodide) → PbI₂(s) (Lead iodide - yellow precipitate) + 2KNO₃(aq) (Potassium nitrate)
      • Neutralization Reaction: Reaction between an acid and a base to form salt and water.
        • HCl(aq) (Acid) + NaOH(aq) (Base) → NaCl(aq) (Salt) + H₂O(l) (Water)

  • e) Oxidation and Reduction (Redox Reactions):

    • Oxidation:
      • Gain of Oxygen
      • Loss of Hydrogen
      • Loss of Electrons (LEO - Loss of Electrons is Oxidation)
      • Increase in oxidation state
    • Reduction:
      • Loss of Oxygen
      • Gain of Hydrogen
      • Gain of Electrons (GER - Gain of Electrons is Reduction)
      • Decrease in oxidation state
    • Redox Reaction: A reaction where oxidation and reduction occur simultaneously.
    • Oxidizing Agent (Oxidant): Substance that gets reduced and causes oxidation of another substance (supplies oxygen or removes hydrogen or accepts electrons). Examples: O₂, Cl₂, KMnO₄, K₂Cr₂O₇, H₂O₂, HNO₃.
    • Reducing Agent (Reductant): Substance that gets oxidized and causes reduction of another substance (removes oxygen or supplies hydrogen or donates electrons). Examples: H₂, C, CO, Metals (Na, Mg, Al), H₂S.
    • Examples:
      • CuO(s) (Black) + H₂(g) --Heat--> Cu(s) (Brown) + H₂O(l)
        • CuO is reduced (loses O) to Cu. CuO is the oxidizing agent.
        • H₂ is oxidized (gains O) to H₂O. H₂ is the reducing agent.
      • ZnO + C → Zn + CO
        • ZnO is reduced to Zn. ZnO is the oxidizing agent.
        • C is oxidized to CO. C is the reducing agent.
      • MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
        • MnO₂ is reduced (loses O) to MnCl₂. MnO₂ is the oxidizing agent.
        • HCl is oxidized (loses H, Cl gains charge from -1 to 0) to Cl₂. HCl is the reducing agent.
      • 2Mg(s) + O₂(g) → 2MgO(s)
        • Mg is oxidized (gains O / loses electrons). Mg is the reducing agent.
        • O₂ is reduced (gains electrons). O₂ is the oxidizing agent.

4. Effects of Oxidation in Everyday Life:

  • a) Corrosion:

    • Definition: The slow degradation of metals by the action of air (oxygen), moisture, or chemicals (like acids) on their surface.
    • Rusting of Iron: Iron reacts with oxygen and moisture (water) to form hydrated ferric oxide (rust).
      • 4Fe(s) + 3O₂(g) + xH₂O(l) → 2Fe₂O₃·xH₂O(s) (Rust - reddish-brown, flaky)
      • Conditions required: Presence of both oxygen and water/water vapour. Impurities, salts (like NaCl), and acids accelerate rusting.
    • Other Examples: Tarnishing of silver (black Ag₂S coating), Green coating on copper (basic copper carbonate - CuCO₃·Cu(OH)₂).
    • Prevention: Painting, Oiling, Greasing, Galvanizing (coating with zinc), Chrome plating, Anodizing (for aluminium), Alloying (making stainless steel).

  • b) Rancidity:

    • Definition: The aerial oxidation of fats and oils present in food materials, resulting in unpleasant smell and taste.
    • Cause: Oxidation of fatty acids and oils.
    • Prevention:
      • Adding Antioxidants: Substances that prevent oxidation (e.g., BHA - Butylated Hydroxyanisole, BHT - Butylated Hydroxytoluene).
      • Packaging in Nitrogen Gas: Flushing bags (like chips packets) with nitrogen (an inert gas) prevents contact with oxygen.
      • Refrigeration: Slows down the oxidation process.
      • Storing in Air-tight Containers: Limits exposure to oxygen.
      • Storing away from Light: Light can accelerate rancidity.

Multiple Choice Questions (MCQs):

  1. Which of the following processes involves a chemical reaction?
    (a) Storing of oxygen gas under pressure in a gas cylinder
    (b) Liquefaction of air
    (c) Keeping petrol in a china dish in the open
    (d) Heating copper wire in the presence of air at high temperature

  2. In the double displacement reaction between aqueous potassium iodide and aqueous lead nitrate, a yellow precipitate of lead iodide is formed. While performing the activity if lead nitrate is not available, which of the following can be used in place of lead nitrate?
    (a) Lead sulphate (insoluble)
    (b) Lead acetate (soluble)
    (c) Ammonium nitrate (soluble)
    (d) Potassium sulphate (soluble)

  3. Which of the following is (are) an endothermic process(es)?
    (i) Dilution of sulphuric acid
    (ii) Sublimation of dry ice
    (iii) Condensation of water vapours
    (iv) Evaporation of water
    (a) (i) and (iii)
    (b) (ii) only
    (c) (iii) only
    (d) (ii) and (iv)

  4. Identify the reducing agent in the following reaction:
    MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
    (a) MnO₂
    (b) HCl
    (c) MnCl₂
    (d) Cl₂

  5. Which of the following represents a balanced chemical equation?
    (a) Fe + Cl₂ → FeCl₃
    (b) 2Mg + O₂ → MgO
    (c) Zn + H₂SO₄ → ZnSO₄ + 2H
    (d) 2Al + 6HCl → 2AlCl₃ + 3H₂

  6. The reaction of H₂ gas with oxygen gas to form water is an example of:
    (a) Combination reaction
    (b) Redox reaction
    (c) Exothermic reaction
    (d) All of the above

  7. What happens when dilute hydrochloric acid is added to iron filings? Choose the correct answer.
    (a) Hydrogen gas and iron(III) chloride are produced.
    (b) Chlorine gas and iron(II) hydroxide are produced.
    (c) No reaction takes place.
    (d) Hydrogen gas and iron(II) chloride are produced.

  8. Which among the following is(are) double displacement reaction(s)?
    (i) Pb + CuCl₂ → PbCl₂ + Cu
    (ii) Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl
    (iii) C + O₂ → CO₂
    (iv) CH₄ + 2O₂ → CO₂ + 2H₂O
    (a) (i) and (iv)
    (b) (ii) only
    (c) (i) and (ii)
    (d) (iii) and (iv)

  9. Silver articles turn black when kept in the open for a few days due to formation of:
    (a) Ag₂O
    (b) Ag₂S
    (c) Ag₂CO₃
    (d) Ag(OH)₂

  10. A dilute ferrous sulphate solution was gradually added to the beaker containing acidified permanganate solution. The light purple colour of the solution fades and finally disappears. Which of the following is the correct explanation for the observation?
    (a) KMnO₄ is an oxidising agent, it oxidises FeSO₄.
    (b) FeSO₄ acts as an oxidising agent and oxidises KMnO₄.
    (c) The colour disappears due to dilution; no reaction is involved.
    (d) KMnO₄ is an unstable compound and decomposes in presence of FeSO₄ to a colourless compound.

Answer Key for MCQs:

  1. (d) - Heating copper forms black copper oxide (CuO), a chemical change. Others are physical changes.
  2. (b) - We need a soluble lead salt to provide Pb²⁺ ions for the reaction with I⁻ ions. Lead acetate is soluble, while lead sulphate is not.
  3. (d) - Sublimation (solid to gas) and evaporation (liquid to gas) both require energy absorption (endothermic). Dilution of acid and condensation are exothermic.
  4. (b) - HCl loses hydrogen (or Cl⁻ is oxidized to Cl₂) and causes the reduction of MnO₂. Therefore, HCl is the reducing agent.
  5. (d) - Check atom count on both sides. Only (d) has equal numbers of Al, H, and Cl atoms on both sides.
  6. (d) - Two substances combine (Combination), H₂ is oxidized and O₂ is reduced (Redox), and heat is released (Exothermic).
  7. (d) - Iron reacts with dilute HCl to form iron(II) chloride (FeCl₂) and hydrogen gas (Fe + 2HCl → FeCl₂ + H₂). Iron(III) chloride (FeCl₃) is typically formed with stronger oxidizing agents or chlorine gas.
  8. (b) - Only reaction (ii) involves the exchange of ions (Na⁺ with Ba²⁺, SO₄²⁻ with Cl⁻). Reaction (i) is displacement. Reactions (iii) and (iv) are combination/combustion.
  9. (b) - Silver reacts with hydrogen sulphide (H₂S) present in the air to form a black coating of silver sulphide (Ag₂S).
  10. (a) - KMnO₄ (potassium permanganate) is a strong oxidizing agent (purple colour). It oxidizes Fe²⁺ (in FeSO₄) to Fe³⁺ and gets reduced itself to a colourless Mn²⁺ compound in acidic medium.

Make sure you understand the reasoning behind each answer. Revise these concepts thoroughly. Good luck with your preparation!

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