NCERT Science Notes - Class 10
Chapter 1 - Chemical Reactions and Equations

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NCERT Science Notes - Class 10
Chapter 1 - Chemical Reactions and Equations

    Introduction

      • Milk left at room temperature during summers:
        • Change: Spoils, sours
        • Nature of change: Chemical (fermentation due to bacterial activity)
      • Iron tawa/pan/nail left exposed to humid atmosphere:
        • Change: Rusting
        • Nature of change: Chemical (oxidation of iron in the presence of moisture)
      • Grapes get fermented:
        • Change: Formation of wine or alcohol
        • Nature of change: Chemical (fermentation due to yeast activity)
      • Food is cooked:
        • Change: Texture, color, taste, and aroma change
        • Nature of change: Both physical and chemical (cooking involves heat-induced chemical changes)
      • Food gets digested in our body:
        • Change: Breakdown of food into simpler substances
        • Nature of change: Chemical (digestive enzymes breaking down complex molecules)
      • We respire:
        • Change: Oxygen intake, carbon dioxide release
        • Nature of change: Chemical (cellular respiration)


      Activity 1.1:

      • Change: Magnesium ribbon burns into magnesium oxide
      • Nature of change: Chemical
      • Observations: Dazzling white flame, formation of magnesium oxide powder


      Activity 1.2:

      • Change: Formation of a yellow precipitate
      • Nature of change: Chemical
      • Observation: Precipitate formation upon mixing lead nitrate and potassium iodide solutions


      Activity 1.3:

      • Change: Zinc reacts with acid
      • Nature of change: Chemical
      • Observations: Gas evolution, change in temperature


      General Observations for Detecting Chemical Reactions:

      • Change in state
      • Change in color
      • Evolution of gas
      • Change in temperature


      Conclusion:

      • Various types of chemical reactions occur around us.
      • Observations like changes in state, color, gas evolution, and temperature help identify chemical reactions.
      • Further study will explore types of chemical reactions and their symbolic representation.

    Chemical Equations

    • Activity 1.1 simplifies chemical reaction descriptions using chemical equations and word-equations.
    • Chemical equations represent reactions concisely.

    Chemical Equations

    Definition:

    • Chemical equations symbolically represent chemical reactions.
    • They have reactants (LHS), products (RHS), and an arrow (→) indicating the reaction direction.

    Word-Equations

    Simplifying Descriptions:

    • Word-equations provide concise descriptions of chemical reactions.
           Example: Magnesium + Oxygen → Magnesium oxide (1.1) (Reactants) (Product)

    Components:

    • Reactants on the LHS, products on the RHS.
    • Arrow (→) shows the reaction direction.
    • Plus signs (+) separate substances.

    Arrow Direction:

    • Arrowhead points to products.

    Balanced Chemical Equations

    1. Introduction
      • The law of conservation of mass states that mass cannot be created or destroyed in a chemical reaction.
      • The number of atoms of each element remains the same before and after a chemical reaction.
      • Chemical equations need to be balanced to ensure mass conservation.
    2. Example Chemical Equation
      • Given word equation: Zinc + Sulphuric acid → Zinc sulphate + Hydrogen
      • Represented as a chemical equation: Zn + H2SO4 → ZnSO4 + H2
      • Number of atoms of different elements in reactants and products are tabulated.
    3. Let's examine the provided chemical equation and the number of atoms of different elements on both sides of the arrow.
      Chemical Equation: Zn + H2SO4 → ZnSO4 + H2

      Here's the analysis of the number of atoms of different elements on both the left-hand side (LHS) and the right-hand side (RHS) of the equation:
      • Zinc (Zn)
        • In the reactants (LHS), there is 1 atom of zinc (Zn).
        • In the products (RHS), there is also 1 atom of zinc (Zn).
      • Hydrogen (H)
        • In the reactants (LHS), there are 2 atoms of hydrogen (H) because there are 2 H atoms in H2SO4.
        • In the products (RHS), there are 2 atoms of hydrogen (H) in H2.
      • Sulphur (S)
        • In the reactants (LHS), there is 1 atom of sulfur (S) in H2SO4.
        • In the products (RHS), there is also 1 atom of sulfur (S) in ZnSO4.
      • Oxygen (O)
        • In the reactants (LHS), there are 4 atoms of oxygen (O) in H2SO4.
        • In the products (RHS), there are 4 atoms of oxygen (O) in ZnSO4.
      As we can see, the number of atoms of each element is the same on both sides of the arrow, which means that this chemical equation - Zn + H2SO4 → ZnSO4 + H2 is already balanced. It follows the law of conservation of mass because the total number of atoms of each element remains the same before and after the chemical reaction.

    Balancing an equation

    Let's start with the process of balancing the chemical equation:

    Chemical Equation: Fe + H2O → Fe3O4 + H2 (1.5)


    Step I: Drawing Boxes 

    To begin balancing the equation, we draw boxes around each formula without changing anything inside them: 

    [ Fe ] + [ H2O ] → [ Fe3O] + [ H]


    Step II: Listing the Number of Atoms 

    Now, let's list the number of atoms of different elements present in the unbalanced equation (1.5) based on your table:

    • Iron (Fe):
      • In the reactants (LHS), there is 1 atom of iron (Fe).
      • In the products (RHS), there are 3 atoms of iron (Fe) in Fe3O4.
    • Hydrogen (H):
      • In the reactants (LHS), there are 2 atoms of hydrogen (H) in H2O.
      • In the products (RHS), there are 2 atoms of hydrogen (H) in H2.
    • Oxygen (O):
      • In the reactants (LHS), there is 1 atom of oxygen (O) in H2O.
      • In the products (RHS), there are 4 atoms of oxygen (O) in Fe3O4.
    As you can see, the number of iron (Fe) and hydrogen (H) atoms is not balanced on both sides of the equation, but the number of oxygen (O) atoms is balanced. We will proceed to balance the remaining elements in the subsequent steps.


    Step III: Balancing Oxygen Atoms

    In Step III, we selected the compound Fe3O4, which contains the element oxygen (O) with the highest number of atoms. Since there are four oxygen atoms on the right-hand side (RHS) and only one on the left-hand side (LHS), we need to balance the oxygen atoms. To do this, we add a coefficient of '4' in front of H2O to equalize the number of oxygen atoms.

    This results in the partly balanced equation: Fe + 4 H2O → Fe3O4 + H2 (1.6) 

    Now, the number of oxygen atoms on both sides is balanced.


    Step IV: Balancing Hydrogen Atoms 

    After balancing the oxygen atoms, we proceed to balance the hydrogen (H) atoms in the partly balanced equation (1.6). To equalize the number of hydrogen atoms, we make the number of molecules of hydrogen (H2) on the right-hand side (RHS) equal to '4.' Here's the updated equation: Fe + 4 H2O → Fe3O4 + 4 H2 (1.7)
    Now, the number of hydrogen atoms on both sides is also balanced. At this point, we have balanced the equation for the number of atoms of each element, including iron (Fe), hydrogen (H), and oxygen (O). The equation is now fully balanced:

    Balanced Equation: Fe + 4 H2O → Fe3O4 + 4 H2 (1.7)

    This balanced equation follows the law of conservation of mass, ensuring that the total number of atoms of each element remains the same before and after the chemical reaction.


    Step V: Balancing Iron (Fe)

    • Examine the equation and identify the unbalanced element, which is iron (Fe).
    • Initially, there's 1 iron atom (Fe) on the left side (LHS) and 3 in Fe3O4 on the right side (RHS).
    • To balance iron, we put a coefficient of '3' in front of Fe on the LHS.
    • The equation becomes: 3 Fe + H2O → Fe3O4 + H2

    Step VI: Checking the Balanced Equation
    • Examine the equation to verify that all elements are balanced.
    • The equation now is: 3 Fe + H2O → Fe3O4 + H2
    • Count the atoms of each element on both sides of the equation.
    • Ensure that the numbers of atoms of elements on both sides are equal.

    Step VII: Adding Physical States (Optional)
    • To provide additional information, add physical states for the reactants and products.
    • The balanced equation becomes: 3 Fe(s) + H2O(g) → Fe3O4(s) + H2(g)
    Now, Eq. (1.2) is balanced and includes the physical states. The coefficients are adjusted to ensure that the law of conservation of mass is satisfied.


    In chemical equations, additional information about the specific conditions under which the reaction occurs, such as temperature, pressure, or the presence of a catalyst, may be provided as annotations above or below the reaction arrow. These conditions can significantly affect the outcome of the chemical reaction and are important to consider when describing the reaction in detail.

    Example 1:                                                   

  1. Chemical Equation: 
                                       
                                                         340 atm     
          CO(g) + 2H2(g)      →     CH3OH(l) (1.11)
  2. In this equation, carbon monoxide (CO) and hydrogen gas (H2) are the reactants, and methanol (CH3OH) is the product.
  3. Above the arrow, the number "340 atm" indicates the pressure at which this chemical reaction is conducted.
  4. This implies that the given chemical reaction occurs at a pressure of 340 atmospheres (atm).

  5. Example 2 (Photosynthesis):
    • Chemical Equation: 


                                                    Sunlight
            6CO(aq) + 12H2O(l)                    C6H12O6(aq) + 6O2(aq) + 6H2O(l) (1.12)                                                             
                                                                          Chlorophyll
    • In this equation representing photosynthesis, carbon dioxide (CO) and water (H2O) are the reactants, and glucose (C6H12O6), oxygen gas (O2), and water are the products.
    • Above the arrow, the word "Sunlight" indicates that sunlight is a necessary reaction condition for photosynthesis to occur.
    • This implies that photosynthesis is a light-dependent process and requires sunlight as a catalyst.

    Types of Chemical Reactions

    • In chemical reactions, atoms of one element don't change into atoms of another element, and no atoms disappear or appear magically.
    • Chemical reactions involve breaking and forming bonds between atoms to create new substances.

    Combination Reaction

    • A combination reaction forms a single product from two or more reactants.
    • Activity 1.4: Calcium oxide (CaO) reacts with water (H2O) to produce slaked lime (Ca(OH)2) and release heat (exothermic).
      • Chemical equation: CaO(s) + H2O(l) → Ca(OH)2(aq) + Heat (1.13).
    Examples of Combination Reactions:
    • (i) Burning of coal:
      • C(s) + O2(g) → CO2(g) (1.15)
    • (ii) Formation of water from hydrogen (H2) and oxygen (O2):
      • 2H2(g) + O2(g) → 2H2O(l) (1.16)
    • Combination reactions occur when two or more substances combine to form a single product.
    Exothermic Reactions:
    • Reactions that release heat along with product formation are called exothermic chemical reactions.
    • Examples include the burning of natural gas and respiration.
    • Respiration involves the breakdown of food into simpler substances to release energy.
      • Example: C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy (1.18)
    • Decomposition of vegetable matter into compost is also exothermic.
    Identifying the Reaction in Activity 1.1:
    • In Activity 1.1, where heat is given out along with the formation of a single product (slaked lime), it is a combination reaction.

    Decomposition Reaction

    Activity 1.5: Decomposition of Ferrous Sulphate
    • Take 2 g of ferrous sulphate crystals in a dry boiling tube.
    • Observe and note the color of the crystals.
    • Heat the boiling tube over a flame.
    • Observe the change in the color of the crystals after heating.
    • Results:
      • The green color of ferrous sulphate crystals changes.
      • A characteristic sulfur smell is noticed.
    • Chemical Reaction:
                                                  
                                                              Heat

                  2FeSO4 (s)     →     Fe2O3 (s)  + SO2 (g) + SO3 (g) (1.19)
                      (Ferrous sulphate)                   (Ferric oxide)

    • Explanation:
      • This is a decomposition reaction where a single substance breaks down into simpler substances.
      • Ferrous sulphate crystals (FeSO4, 7H2O) lose water when heated, causing a color change.
      • The decomposition results in ferric oxide (Fe2O3), sulphur dioxide (SO2), and sulphur trioxide (SO3). Ferric oxide is solid, while SO2 and SO3 are gases.
    Thermal Decomposition:
    • When decomposition occurs due to heating, it's called thermal decomposition.
    • Example: Decomposition of calcium carbonate (CaCO3) into calcium oxide (CaO) and carbon dioxide (CO2).
      • CaCO3 (s) (Limestone) → Heat → CaO (s) (Quick lime) + CO2 (g) (1.20)

    Activity 1.6: Decomposition of Lead Nitrate
    • Take 2 g of lead nitrate powder in a boiling tube.
    • Heat it over a flame.
    • Observe and note any changes.
    • Results:
      • Brown fumes, identified as nitrogen dioxide (NO2), are emitted.
    • Chemical Reaction:

                                                         Heat
            2Pb(NO3)2 (s)    →       2PbO (s)     +    4NO2 (g)     +    O2 (g)  (1.21)
                      (Lead nitrate)                             (Lead oxide)               (Nitrogen dioxide)              (Oxygen)
    • Explanation:
      • This is another example of a decomposition reaction.
      • Lead nitrate breaks down upon heating, producing brown nitrogen dioxide fumes.
    These activities demonstrate how substances can break down into simpler components through decomposition reactions, either due to heating or other factors.

    Activity 1.7 - Electrolysis of Water

    Setup:

    • Plastic mug with two holes at the base, fitted with rubber stoppers.
    • Carbon electrodes inserted into the rubber stoppers.
    • 6-volt battery connected to the carbon electrodes.
    • Mug filled with water, with electrodes immersed.
    • A few drops of dilute sulphuric acid added to the water.
    • Two test tubes filled with water inverted over the carbon electrodes.

    Procedure and Observations:

    1. Switch on the current and leave the setup undisturbed.
    2. Bubbles form at both electrodes.
    3. The volume of gas collected in both test tubes may not be the same.
    4. Carefully remove the test tubes once they are filled with gases.
    5. Test the gases by bringing a burning candle close to the mouth of each test tube.

    Results:

    • One test tube collects hydrogen gas (H2), and the other collects oxygen gas (O2).
    • The volume of hydrogen gas is usually double that of oxygen gas.


    Activity 1.8 - Decomposition of Silver Chloride in Sunlight

    Procedure and Observations:

    1. Take 2 g of white silver chloride (AgCl) in a china dish.
    2. Place the china dish in sunlight for some time.
    3. Observe the color of silver chloride after exposure to sunlight.
    4. Note the color change.

    Results:

    • White silver chloride turns grey in sunlight.
    • This is due to the decomposition of silver chloride into silver (Ag) and chlorine gas (Cl2) when exposed to light.

    Additional Information:

    • Silver bromide (AgBr) behaves similarly, decomposing into silver and bromine gas (Br2) in sunlight.
    • These reactions are utilized in black and white photography.
    • The decomposition reactions are driven by the energy from light.

    Barium Hydroxide and Ammonium Chloride Reaction

    • Mixing 2 g of barium hydroxide (Ba(OH)2) with 1 g of ammonium chloride (NH4Cl) in a test tube.
    • Feeling a cooling sensation when touching the bottom of the test tube.
    • This is an endothermic reaction, as it absorbs heat from the surroundings.

    Displacement Reaction

    Activity 1.9 : Procedure

    1. Take three iron nails and clean them using sandpaper.
    2. Prepare two test tubes labeled as (A) and (B).
    3. Fill each test tube with approximately 10 mL of copper sulphate solution.
    4. Tie two of the iron nails together using a thread.
    5. Immerse the tied iron nails carefully in the copper sulphate solution in test tube B for about 20 minutes.
    6. Keep one iron nail aside for comparison.
    7. After 20 minutes, remove the iron nails from the copper sulphate solution.

    Observations:

    • Compare the intensity of the blue color of the copper sulphate solutions in test tubes (A) and (B).
    • Also, compare the color of the iron nails dipped in the copper sulphate solution with the one kept aside.

    Results and Explanation:

    • The iron nail becomes brownish in color.
    • The blue color of copper sulphate solution in test tube (B) fades.

    Chemical Reaction:

            Fe(s)      +      CuSO4 (aq)       →      FeSO4 (aq)       +      Cu(s)                 (1.24)
                                           (Copper sulphate)                    (Iron sulphate)                   (Copper)


    Explanation:

    • In this reaction, iron displaces copper from the copper sulphate solution.
    • Such a reaction is known as a displacement reaction.
    • Other examples of displacement reactions include:
      • Zn(s) + CuSO4 (aq)  → ZnSO4 (aq)  + Cu(s)  (1.25)
                                    (Copper sulphate)           (Zinc sulphate)          (Copper)

      • Pb(s) + CuCl2 (aq) → PbCl2 (aq)  +  Cu(s)  (1.26)
                                   (Copper chloride)       (Lead chloride)        (Copper)
    • Zinc and lead are more reactive elements than copper, so they can displace copper from its compounds.

    Double Displacement Reaction

    Activity 1.10 : Procedure
    1. Take about 3 mL of sodium sulphate solution in one test tube.
    2. In another test tube, take about 3 mL of barium chloride solution.
    3. Mix the two solutions.
    Observation:
    • A white substance that is insoluble in water is formed.
    Result:
    • The insoluble substance formed is known as a precipitate.
    • Reactions that produce a precipitate are called precipitation reactions.
    Chemical Reaction: Na 2SO 4 (aq) + BaCl 2 (aq) → BaSO 4 (s) + 2NaCl(aq) (1.27)

    Explanation:

    • In this reaction, sodium sulphate (Na2SO4) reacts with barium chloride (BaCl2) to produce barium sulphate (BaSO4) as a white precipitate and sodium chloride (NaCl) which remains in the solution.
    • This type of reaction, where there is an exchange of ions between the reactants, is known as a double displacement reaction.

    Oxidation and Reduction

    Activity 1.11 : Observation

    • When a china dish containing about 1 g of copper powder is heated, the surface of the copper powder becomes coated with black copper(II) oxide (CuO).
    Explanation:
    • The black substance, copper(II) oxide (CuO), forms because oxygen is added to copper during heating.
    • This reaction can be represented as: 2Cu + O2 → Heat → 2CuO (1.28)
    • If hydrogen gas (H2) is passed over the heated copper oxide (CuO), the black coating on the surface turns brown. This is because the reverse reaction takes place, and copper is obtained from copper oxide.
    • The reaction with hydrogen can be represented as: CuO + H2 → Heat → Cu + H2O (1.29)
    Oxidation and Reduction:
    • If a substance gains oxygen during a reaction, it is said to be oxidized.
    • If a substance loses oxygen during a reaction, it is said to be reduced.
    • In reaction (1.29), copper(II) oxide is losing oxygen and is being reduced, while hydrogen is gaining oxygen and is being oxidized.
    • Reactions in which one reactant gets oxidized while the other gets reduced are called oxidation-reduction reactions or redox reactions.
    Examples of Redox Reactions:
    1. ZnO + C → Zn + CO (1.31) - Carbon is oxidized to CO, and ZnO is reduced to Zn.
    2. MnO2 + 4HCl → MnCl2 + 2H2O + Cl2 (1.32) - HCl is oxidized to Cl2, and MnO2 is reduced to MnCl2.
    General Rule:
    • If a substance gains oxygen or loses hydrogen during a reaction, it is oxidized.
    • If a substance loses oxygen or gains hydrogen during a reaction, it is reduced.
    Recalling Activity 1.1:
    • In Activity 1.1, where a magnesium ribbon burns with a dazzling flame in air (oxygen) and changes into a white substance, magnesium oxide (MgO), magnesium is oxidized in this reaction.

    Corrosion

    • Corrosion is a common phenomenon where metals, especially iron, tend to deteriorate and form undesirable compounds when exposed to substances like moisture, acids, etc.
    • For example, iron articles are shiny when new but gradually get coated with a reddish-brown powder when left for some time, a process known as rusting.
    • Other metals like copper and silver can also get tarnished, leading to the formation of coatings on their surfaces. Silver forms a black coating, while copper develops a greenish one.
    • Corrosion is a major issue as it damages various objects made of metals, including car bodies, bridges, iron railings, and ships, primarily those composed of iron.
    • The problem of corrosion is so significant that substantial amounts of money are spent annually on replacing damaged iron objects.
    • Further details about corrosion are discussed in Chapter 3.

    Rancidity

    • Rancidity refers to the unpleasant taste and smell that develops in food materials containing fats and oils when they undergo oxidation over time.
    • Foods containing fats and oils are particularly susceptible to rancidity, and this change in taste and smell can make the food unappetizing.
    • To prevent oxidation, which leads to rancidity, substances known as antioxidants are often added to such foods.
    • Storing food in airtight containers can also help slow down the oxidation process.
    • Notably, manufacturers of products like chips use techniques such as flushing bags with gases like nitrogen to prevent the chips from getting oxidized and becoming rancid.

    CBSE Class 10 Science Chapter 1 - Chemical Reactions and Equations 

    Class 10 CBSE Important Questions and Answers Chapter 1 - Chemical Reactions and Equations


    Class 10 NCERT Chapter 1 - Chemical Reactions and Equations AJs Chalo Seekhen Class 10 CBSE Important Questions and Answers Chapter 1 - Chemical Reactions and Equations ajs notes history chapter 1  ajs class 10 chapter 1 imp questions

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