Class 10 Notes

Class 10 Science Notes Chapter 1 (Chemical reactions and equations) – Science

Philoid

Philoid

— 6 min read

Science

Okay, let's begin with the detailed notes for Chapter 1: Chemical Reactions and Equations, focusing on aspects relevant for government exam preparation based on the NCERT Class 10 syllabus.

Chapter 1: Chemical Reactions and Equations

1. Chemical Reactions:

  • Definition: A process involving the rearrangement of the structure of molecules or ions into new substances with new physical and chemical properties. Essentially, old bonds are broken, and new bonds are formed.
  • Chemical Change vs. Physical Change:
    • Physical Change: Alters the form or appearance of a substance but not its chemical composition (e.g., melting ice, boiling water, cutting paper). Usually reversible.
    • Chemical Change: Results in the formation of one or more new substances with different properties (e.g., burning wood, rusting of iron, digestion of food). Usually irreversible or difficult to reverse.
  • Identifying Chemical Reactions: Observations that suggest a chemical reaction has occurred:
    • Change in State: Reactants might be solid/liquid/gas, and products might be in a different state. (e.g., Burning candle wax (solid) produces CO₂ (gas) and H₂O (gas/liquid)).
    • Change in Colour: Formation of new substances often leads to colour changes. (e.g., Iron reacting with copper sulphate solution changes the blue solution to greenish).
    • Evolution of a Gas: Bubbles indicate gas formation. (e.g., Zinc reacting with dilute sulphuric acid produces hydrogen gas).
    • Change in Temperature:
      • Exothermic Reaction: Releases heat, causing the temperature of the surroundings to rise. (e.g., Reaction of quicklime (CaO) with water).
      • Endothermic Reaction: Absorbs heat, causing the temperature of the surroundings to fall. (e.g., Reaction of barium hydroxide with ammonium chloride).
    • Formation of a Precipitate: Formation of an insoluble solid substance. (e.g., Reaction of lead nitrate with potassium iodide forms a yellow precipitate of lead iodide).

2. Chemical Equations:

  • Definition: A symbolic representation of a chemical reaction using chemical formulas and symbols.
  • Components:
    • Reactants: Substances that undergo chemical change. Written on the left-hand side (LHS).
    • Products: New substances formed during the reaction. Written on the right-hand side (RHS).
    • Arrow (→): Separates reactants from products and indicates the direction of the reaction ('yields' or 'produces').
    • Plus Sign (+): Used between multiple reactants or products.
  • Skeletal Chemical Equation: An unbalanced chemical equation that only represents the identities of reactants and products using their formulas. (e.g., Mg + O₂ → MgO).
  • Balanced Chemical Equation: An equation where the number of atoms of each element is equal on both the reactant (LHS) and product (RHS) sides. This adheres to the Law of Conservation of Mass.
    • Law of Conservation of Mass: Mass can neither be created nor destroyed in a chemical reaction. The total mass of reactants must equal the total mass of products. This implies the number of atoms of each element must remain the same before and after the reaction.
  • Writing and Balancing Chemical Equations:
    • Step 1: Write the skeletal equation using correct chemical formulas.
    • Step 2: List the number of atoms of each element on LHS and RHS.
    • Step 3: Start balancing with the element that appears in the maximum number of compounds or has the maximum number of atoms (often polyatomic ions can be balanced as a unit if they appear unchanged on both sides). Balance metals and non-metals first. Balance hydrogen and oxygen last.
    • Step 4: Use coefficients (numbers placed before formulas) to equalize the atom count. Never change the chemical formula of a substance.
    • Step 5: Verify that the number of atoms for each element is equal on both sides.
    • Step 6 (Optional but good practice): Write the physical states of reactants and products: (s) for solid, (l) for liquid, (g) for gas, and (aq) for aqueous solution (dissolved in water). Reaction conditions like temperature, pressure, or catalyst can be written above or below the arrow.
    • Example: Balancing Fe + H₂O → Fe₃O₄ + H₂
      1. Skeletal: Fe + H₂O → Fe₃O₄ + H₂
      2. Atoms: LHS (Fe=1, H=2, O=1), RHS (Fe=3, H=2, O=4)
      3. Balance Fe: 3Fe + H₂O → Fe₃O₄ + H₂ (LHS: Fe=3, H=2, O=1; RHS: Fe=3, H=2, O=4)
      4. Balance O: 3Fe + 4H₂O → Fe₃O₄ + H₂ (LHS: Fe=3, H=8, O=4; RHS: Fe=3, H=2, O=4)
      5. Balance H: 3Fe + 4H₂O → Fe₃O₄ + 4H₂ (LHS: Fe=3, H=8, O=4; RHS: Fe=3, H=8, O=4) - Balanced.
      6. With states (assuming steam is used): 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

3. Types of Chemical Reactions:

  • a) Combination Reaction:
    • Definition: Two or more reactants combine to form a single product. (A + B → AB)
    • Examples:
      • Burning of coal: C(s) + O₂(g) → CO₂(g)
      • Formation of water: 2H₂(g) + O₂(g) → 2H₂O(l)
      • Quicklime (Calcium oxide) reacting with water: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat (Slaked lime - This is highly exothermic)
  • b) Decomposition Reaction:
    • Definition: A single reactant breaks down into two or more simpler products. Requires energy input (heat, light, or electricity). (AB → A + B)
    • Types based on energy source:
      • Thermal Decomposition: Decomposition by heat.
        • Ferrous sulphate crystals: 2FeSO₄(s) --(Heat)→ Fe₂O₃(s) + SO₂(g) + SO₃(g) (Green crystals change to brown solid; smell of burning sulphur)
        • Calcium carbonate (limestone): CaCO₃(s) --(Heat)→ CaO(s) + CO₂(g) (Used in cement industry)
        • Lead nitrate: 2Pb(NO₃)₂(s) --(Heat)→ 2PbO(s) + 4NO₂(g) + O₂(g) (White solid changes to yellow solid; brown fumes of NO₂ evolved)
      • Electrolytic Decomposition (Electrolysis): Decomposition by passing electric current.
        • Water: 2H₂O(l) --(Electric Current)→ 2H₂(g) + O₂(g) (H₂ collects at cathode, O₂ at anode; volume of H₂ is double that of O₂)
      • Photolytic Decomposition (Photolysis): Decomposition by light energy.
        • Silver chloride: 2AgCl(s) --(Sunlight)→ 2Ag(s) + Cl₂(g) (White solid turns grey)
        • Silver bromide: 2AgBr(s) --(Sunlight)→ 2Ag(s) + Br₂(g) (Used in black and white photography)
  • c) Displacement Reaction:
    • Definition: A more reactive element displaces a less reactive element from its compound (salt solution). (A + BC → AC + B, where A is more reactive than B). Reactivity is determined by the reactivity series.
    • Examples:
      • Iron nail in copper sulphate solution: Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) (Blue solution turns pale green; reddish-brown copper deposits on iron)
      • Zinc in copper sulphate solution: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s) (Blue solution turns colourless; reddish-brown copper deposits)
      • Lead in copper chloride solution: Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)
  • d) Double Displacement Reaction:
    • Definition: Two compounds react by exchanging their ions to form two new compounds. (AB + CD → AD + CB)
    • Often involves the formation of a precipitate (insoluble solid). These are also called Precipitation Reactions.
    • Examples:
      • Sodium sulphate and Barium chloride: Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq) (White precipitate of Barium sulphate formed)
      • Lead nitrate and Potassium iodide: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq) (Yellow precipitate of Lead iodide formed)
    • Neutralization Reaction: Reaction between an acid and a base to form salt and water. This is also a type of double displacement reaction.
      • NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)
  • e) Oxidation and Reduction (Redox Reactions):
    • Oxidation:
      • Gain of Oxygen
      • Loss of Hydrogen
      • (Higher classes: Loss of electrons)
    • Reduction:
      • Loss of Oxygen
      • Gain of Hydrogen
      • (Higher classes: Gain of electrons)
    • Redox Reaction: A reaction where oxidation and reduction occur simultaneously.
    • Oxidizing Agent: The substance that causes oxidation (by providing oxygen or removing hydrogen) and gets reduced itself.
    • Reducing Agent: The substance that causes reduction (by removing oxygen or providing hydrogen) and gets oxidized itself.
    • Examples:
      • CuO(s) + H₂(g) --(Heat)→ Cu(s) + H₂O(l)
        • CuO loses oxygen → Reduced (CuO is the oxidizing agent)
        • H₂ gains oxygen → Oxidized (H₂ is the reducing agent)
      • ZnO(s) + C(s) → Zn(s) + CO(g)
        • ZnO loses oxygen → Reduced (ZnO is the oxidizing agent)
        • C gains oxygen → Oxidized (C is the reducing agent)
      • MnO₂(s) + 4HCl(aq) → MnCl₂(aq) + 2H₂O(l) + Cl₂(g)
        • MnO₂ loses oxygen → Reduced (MnO₂ is the oxidizing agent)
        • HCl loses hydrogen (to form Cl₂) → Oxidized (HCl is the reducing agent)

4. Effects of Oxidation in Everyday Life:

  • a) Corrosion:
    • Definition: The slow degradation (eating away) of metals when exposed to substances like moisture, air (oxygen), acids, etc.
    • Examples:
      • Rusting of Iron: Formation of hydrated iron(III) oxide (Fe₂O₃.xH₂O) - reddish-brown coating. Requires both oxygen and water.
      • Tarnishing of Silver: Formation of black silver sulphide (Ag₂S) due to reaction with H₂S in the air.
      • Green coating on Copper: Formation of basic copper carbonate (CuCO₃.Cu(OH)₂) upon exposure to moist air (CO₂, O₂, H₂O).
    • Prevention: Painting, oiling, greasing, galvanizing (coating iron with zinc), chromium plating, anodizing (forming a thick oxide layer, e.g., on aluminium), alloying (mixing metals).
  • b) Rancidity:
    • Definition: The oxidation of fats and oils present in food materials, resulting in unpleasant smell and taste, making the food unfit for consumption.
    • Prevention:
      • Adding Antioxidants: Substances that prevent oxidation (e.g., BHA - Butylated Hydroxyanisole, BHT - Butylated Hydroxytoluene).
      • Vacuum Packing: Removing air from packaging.
      • Flushing with Nitrogen: Replacing air (oxygen) with less reactive nitrogen gas (e.g., in chips packets).
      • Refrigeration: Slows down the rate of oxidation.
      • Storing food in airtight containers.

Key Focus Areas for Government Exams:

  1. Balancing Chemical Equations: Expect questions requiring balancing or identifying correctly balanced equations.
  2. Identifying Types of Reactions: Given a reaction, identify its type (Combination, Decomposition, Displacement, Double Displacement, Redox).
  3. Specific Examples: Remember key examples, especially colour changes, precipitate formation (colours of precipitates like BaSO₄, PbI₂), gases evolved (H₂, O₂, CO₂, SO₂, NO₂ - brown fumes), and decomposition reactions (thermal, electrolytic, photolytic) with their specific products and conditions.
  4. Redox Reactions: Identify substances oxidized, reduced, oxidizing agents, and reducing agents in a given reaction.
  5. Corrosion and Rancidity: Definitions, common examples (rust formula), and methods of prevention.
  6. Law of Conservation of Mass: Understand its implication for balancing equations.
  7. Exothermic vs. Endothermic: Identify reactions based on heat release/absorption (e.g., CaO + H₂O is exothermic, decomposition reactions are generally endothermic).
  8. Reactivity Series (Implicit): Understand the basis for displacement reactions (a more reactive metal displaces a less reactive one).

These notes cover the essential concepts from NCERT Chapter 1. Consistent revision and practice with example questions will be beneficial for exam preparation.

Read more