Gonit Vigyan

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Chemical Reactions and Their Indicators

Introduction

In our daily lives, we come across various situations where substances undergo a change in their identity. These changes can be broadly classified as physical changes and chemical changes. When a chemical change occurs, a new substance with different properties is formed, indicating that a chemical reaction has taken place.

Examples of Chemical Changes in Daily Life

1. Milk left at room temperature during summers – The milk undergoes souring due to bacterial action, forming lactic acid.
2. Iron tawa/pan/nail exposed to humid atmosphere – The iron reacts with oxygen and water to form rust (iron oxide).
3. Fermentation of grapes – The sugars in grapes get converted into alcohol and carbon dioxide due to microbial action.
4. Cooking of food – Heat causes complex reactions that change the taste, texture, and composition of food.
5. Digestion of food – Enzymes break down food into simpler molecules in our body.
6. Respiration – Oxygen reacts with glucose in cells to produce energy, carbon dioxide, and water.

In all the above cases, the nature and identity of the original substance change, confirming the occurrence of a chemical reaction.

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Understanding Chemical Reactions

Activity 1.1 – Burning of Magnesium Ribbon

Objective: To observe the chemical change during the burning of magnesium in air.

Procedure:

• Take a magnesium ribbon (3-4 cm long).
• Clean it using sandpaper to remove the oxide layer.
• Hold it with a pair of tongs and burn it using a spirit lamp or burner.
• Collect the white ash formed in a watch-glass.

Observations:

• The magnesium ribbon burns with a dazzling white flame.
• A white powder (magnesium oxide, MgO) is formed.

Conclusion:

• Magnesium reacts with oxygen in the air to form magnesium oxide (MgO).
• This is an example of a chemical reaction, where a new substance is formed.

Chemical Equation:

$$2Mg + O_2 \rightarrow 2MgO$$

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Activity 1.2 – Reaction of Zinc with Acid

Objective: To observe the reaction between zinc and dilute sulphuric acid.

Procedure:

• Take zinc granules in a test tube or conical flask.
• Add dilute hydrochloric acid (HCl) or sulphuric acid (H₂SO₄).
• Observe the reaction and touch the flask.

Observations:

• Bubbles (hydrogen gas) evolve, indicating a gas is formed.
• The test tube becomes warm, showing a rise in temperature.

Conclusion:

• Zinc reacts with acid, producing hydrogen gas and zinc chloride/sulfate.
• This is an example of a chemical reaction with gas evolution and temperature change.

Chemical Equation:

$$Zn + 2HCl \rightarrow ZnCl_2 + H_2 \uparrow$$

or

$$Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2 \uparrow$$

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Activity 1.3 – Reaction of Lead Nitrate with Potassium Iodide

Objective: To observe a precipitation reaction.

Procedure:

• Take lead nitrate solution (Pb(NO₃)₂) in a test tube.
• Add potassium iodide solution (KI) to it.

Observations:

• A yellow precipitate of lead iodide (PbI₂) is formed.

Conclusion:

• A precipitate formation indicates a chemical reaction.

Chemical Equation:

$$Pb(NO_3)_2 + 2KI \rightarrow PbI_2 \downarrow + 2KNO_3$$

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Indicators of a Chemical Reaction

From the above activities, we can identify a chemical reaction by the following observations:

1. Change in State – Solid, liquid, or gas formation.
2. Change in Colour – Example: Yellow precipitate in lead nitrate reaction.
3. Evolution of Gas – Example: Hydrogen gas in zinc-acid reaction.
4. Change in Temperature – Heat absorption or release, making the container hot or cold.

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Conclusion

• A chemical reaction involves the formation of one or more new substances with different properties from the reactants.
• Various observable changes, such as colour change, gas evolution, precipitation, and temperature change, indicate the occurrence of a chemical reaction.
• These reactions play an essential role in natural and industrial processes.

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1.1 CHEMICAL EQUATIONS

Introduction to Chemical Equations

A chemical equation is a shorthand way to represent a chemical reaction using symbols and formulas. It provides a clear and concise way to express the transformation of reactants into products.

Word Equation Representation

Consider the reaction from Activity 1.1, where magnesium burns in oxygen to form magnesium oxide. The reaction can be written as:

$$\text{Magnesium} + \text{Oxygen} \rightarrow \text{Magnesium Oxide}$$

• Reactants: Magnesium (Mg) and Oxygen (O₂) (Substances undergoing a chemical change)
• Product: Magnesium Oxide (MgO) (New substance formed)

Structure of a Word Equation

• Reactants are written on the left-hand side (LHS).
• Products are written on the right-hand side (RHS).
• The arrow (→) represents the direction of the reaction, showing that reactants change into products.
• A plus sign (+) is used to separate multiple reactants or products.

Symbolic Representation of a Chemical Equation

Instead of writing a word equation, a chemical reaction can be expressed using chemical symbols and formulas.

The reaction:

$$\text{Magnesium} + \text{Oxygen} \rightarrow \text{Magnesium Oxide}$$

can be written as:

$$\mathbf{Mg} + \mathbf{O_2} \rightarrow \mathbf{MgO}$$

This is a more precise and concise representation of the reaction.

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Conclusion

• A chemical equation is a way to represent a chemical reaction in a symbolic form.
• It consists of reactants, an arrow indicating the reaction direction, and products.
• A word equation describes the reaction in words, while a symbolic equation uses chemical formulas.

Here's a detailed note on Writing a Chemical Equation based on the given text:

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1.1.1 Writing a Chemical Equation

Introduction

A chemical reaction can be represented in different ways. While word equations are useful, they can be made shorter and more precise by using chemical symbols and formulas. This representation is called a chemical equation.

Chemical Formula Representation

Instead of writing:

$$\text{Magnesium} + \text{Oxygen} \rightarrow \text{Magnesium Oxide}$$

We use chemical formulas:

$$\mathbf{Mg} + \mathbf{O_2} \rightarrow \mathbf{MgO}$$

This is a more concise and useful way of writing chemical reactions.

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Skeletal Chemical Equations

A skeletal chemical equation is an unbalanced equation that represents the reactants and products but does not ensure that the number of atoms is equal on both sides.

Let’s analyze equation (1.2):

$$\mathbf{Mg} + \mathbf{O_2} \rightarrow \mathbf{MgO}$$

• Left-hand side (LHS):

  • Mg = 1 atom
  • O = 2 atoms (O₂ molecule)

• Right-hand side (RHS):

  • Mg = 1 atom
  • O = 1 atom (in MgO)

Here, oxygen atoms are not equal on both sides:

• LHS: 2 oxygen atoms
• RHS: 1 oxygen atom

Since the number of atoms on both sides is not the same, this equation is unbalanced and is called a skeletal chemical equation.

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Conclusion

• A chemical equation is a symbolic representation of a chemical reaction using chemical formulas.
• A skeletal equation is an unbalanced equation where the number of atoms of each element is not the same on both sides.
• To make a chemical equation correct, we must balance it, ensuring the Law of Conservation of Mass is followed.

1.1.2 Balanced Chemical Equations

Law of Conservation of Mass

Before balancing chemical equations, recall the Law of Conservation of Mass:

"Mass can neither be created nor destroyed in a chemical reaction."

This means the total mass of reactants must be equal to the total mass of products. Therefore, the number of atoms of each element must remain the same before and after a chemical reaction.

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Balancing a Chemical Equation

Let’s analyze the reaction:

$$\text{Zinc} + \text{Sulphuric Acid} \rightarrow \text{Zinc Sulphate} + \text{Hydrogen}$$

Its chemical representation is:

$$\mathbf{Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2}$$

Now, count the number of atoms of each element on both sides:

Element	Reactants (LHS)	Products (RHS)
Zn (Zinc)	1	1
H (Hydrogen)	2	2
S (Sulphur)	1	1
O (Oxygen)	4	4

Since the number of atoms is the same on both sides, the equation is already balanced.

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Balancing Fe + H₂O → Fe₃O₄ + H₂

Now, let’s balance the skeletal equation:

$$\mathbf{Fe + H_2O \rightarrow Fe_3O_4 + H_2}$$

Step-by-Step Balancing

Step 1: Identify the number of atoms in the unbalanced equation.

Element	Reactants (LHS)	Products (RHS)
Fe (Iron)	1	3
H (Hydrogen)	2	2
O (Oxygen)	1	4

Step 2: Balance the oxygen atoms.

• Oxygen appears 4 times in Fe₃O₄ but only once in H₂O.
• So, multiply H₂O by 4:

$$Fe + 4H_2O \rightarrow Fe_3O_4 + H_2$$

Now, oxygen is balanced.

Step 3: Balance hydrogen atoms.

• On the LHS: 4 × H₂O = 8 hydrogen atoms
• On the RHS: H₂ (2 hydrogen atoms)
• To balance hydrogen, multiply H₂ by 4:

$$Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$$

Step 4: Balance iron atoms.

• On the LHS: Only 1 Fe atom
• On the RHS: 3 Fe atoms in Fe₃O₄
• Multiply Fe by 3 on the LHS:

$$\mathbf{3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2}$$

Now, the equation is balanced.

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Step 5: Indicating Physical States

To make the equation more informative, add physical states:

$$\mathbf{3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)}$$

• (s) → Solid
• (l) → Liquid
• (g) → Gas
• (aq) → Aqueous (dissolved in water)

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General Rules for Balancing Equations

1. Write the correct formulas for all reactants and products.
2. Do not change the chemical formulas of any substances.
3. Start balancing elements that appear in complex compounds first.
4. Use the smallest whole numbers as coefficients.
5. Check the final equation by counting atoms of each element.

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Examples of Balanced Equations

1. Formation of Methanol:

$$$$

2. Photosynthesis Reaction:

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Conclusion

• Balancing chemical equations ensures that the Law of Conservation of Mass is followed.
• The hit-and-trial method is used to adjust coefficients.
• Physical states and reaction conditions provide additional information about reactions.

Questions And Answer 

1. Why should a magnesium ribbon be cleaned before burning in air?

Magnesium ribbon should be cleaned before burning to remove the oxide layer (MgO) formed due to its reaction with atmospheric oxygen. This oxide layer can interfere with the burning process and slow down the reaction between magnesium and oxygen. Cleaning ensures that pure magnesium reacts efficiently with oxygen to form magnesium oxide (MgO).

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2. Write the balanced equation for the following chemical reactions:

(i) Hydrogen + Chlorine → Hydrogen chloride

$$H_2 + Cl_2 \rightarrow 2HCl$$

(ii) Barium chloride + Aluminium sulphate → Barium sulphate + Aluminium chloride

$$3BaCl_2 + Al_2(SO_4)_3 \rightarrow 3BaSO_4 + 2AlCl_3$$

(iii) Sodium + Water → Sodium hydroxide + Hydrogen

$$2Na + 2H_2O \rightarrow 2NaOH + H_2$$

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3. Write a balanced chemical equation with state symbols:

(i) Solutions of barium chloride and sodium sulphate in water react to give insoluble barium sulphate and the solution of sodium chloride.

$$BaCl_2 (aq) + Na_2SO_4 (aq) \rightarrow BaSO_4 (s) + 2NaCl (aq)$$

(ii) Sodium hydroxide solution (in water) reacts with hydrochloric acid solution (in water) to produce sodium chloride solution and water.

$$NaOH (aq) + HCl (aq) \rightarrow NaCl (aq) + H_2O (l)$$

Here are detailed notes on Types of Chemical Reactions: Combination Reactions from your provided text:

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1.2 TYPES OF CHEMICAL REACTIONS

1.2.1 Combination Reactions

Definition:

A combination reaction is a type of chemical reaction in which two or more reactants (elements or compounds) combine to form a single product.

Example 1: Reaction of Calcium Oxide with Water

• Activity 1.4:
  • Take a small amount of calcium oxide (quick lime) in a beaker.
  • Slowly add water to it.
  • Touch the beaker and observe the temperature change.
• Observation: The beaker becomes hot.
• Reaction: $$CaO(s) + H_2O(l) → Ca(OH)_2 (aq) + \text{Heat}$$
  • Reactants: Calcium oxide (CaO) and water (H₂O).
  • Product: Calcium hydroxide (Ca(OH)₂), also called slaked lime.
  • Heat is released, making the reaction exothermic.

Example 2: Burning of Coal

$$C(s) + O_2(g) → CO_2(g)$$

• Carbon (C) combines with oxygen (O₂) to form carbon dioxide (CO₂).

Example 3: Formation of Water

$$2H_2(g) + O_2(g) → 2H_2O(l)$$

• Hydrogen (H₂) and oxygen (O₂) combine to form water (H₂O).

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Exothermic Chemical Reactions

Definition:

Exothermic reactions are chemical reactions in which heat is released along with the formation of products.

Examples of Exothermic Reactions:

(i) Burning of Natural Gas

$$CH_4(g) + 2O_2(g) → CO_2(g) + 2H_2O(g) + \text{Heat}$$

• Methane (CH₄) burns in oxygen (O₂) to form carbon dioxide (CO₂) and water (H₂O) with heat release.

(ii) Respiration - An Exothermic Process

$$C_6H_{12}O_6(aq) + 6O_2(aq) → 6CO_2(aq) + 6H_2O(l) + \text{Energy}$$

• Glucose (C₆H₁₂O₆) in our body reacts with oxygen (O₂) to release energy.
• This is how our body generates energy from food.

(iii) Decomposition of Vegetable Matter into Compost

• When vegetable matter decomposes, heat is released, making it an exothermic process.

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Whitewashing Reaction (Limewater Reaction)

• Step 1: Slaked lime (Ca(OH)₂) is produced by mixing quick lime (CaO) with water.
• Step 2: When slaked lime is applied on walls for whitewashing, it reacts with carbon dioxide (CO₂) in air and forms a thin layer of calcium carbonate (CaCO₃).

$$Ca(OH)_2 (aq) + CO_2 (g) → CaCO_3 (s) + H_2O(l)$$

• Calcium carbonate gives a shiny white finish to walls.
• Interesting Fact: The chemical formula for marble is also CaCO₃.

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Key Takeaways:

1. Combination Reactions involve two or more reactants forming a single product.
2. Exothermic Reactions release heat energy.
3. Common examples include burning of coal, respiration, and whitewashing.
4. Whitewashing is an application of combination reactions where slaked lime reacts with CO₂ to form calcium carbonate.

Here are detailed notes on 1.2.2 Decomposition Reactions from your provided text:

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1.2.2 Decomposition Reactions

Definition:

A decomposition reaction is a type of chemical reaction in which a single reactant breaks down into two or more simpler products.

• These reactions usually require an input of energy in the form of heat, light, or electricity to occur.
• Decomposition reactions are generally endothermic, meaning they absorb energy.

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Types of Decomposition Reactions:

1. Thermal Decomposition (Decomposition by Heat)

• Definition: When a compound breaks down into simpler substances upon heating, it is called thermal decomposition.

Example 1: Decomposition of Ferrous Sulphate (Activity 1.5)

• Experiment:
  • Take ferrous sulphate crystals (FeSO₄·7H₂O) in a dry boiling tube.
  • Heat it over a flame.
  • Observation:
    • The green colour of ferrous sulphate turns brown.
    • A pungent smell of burning sulphur is observed.
  • Reaction: $$2FeSO_4(s) \xrightarrow{\text{Heat}} Fe_2O_3(s) + SO_2(g) + SO_3(g)$$
    • Ferrous sulphate (FeSO₄) decomposes to form ferric oxide (Fe₂O₃), sulphur dioxide (SO₂), and sulphur trioxide (SO₃).
    • Ferric oxide (Fe₂O₃) is a brown solid.
    • SO₂ and SO₃ are gases.

Example 2: Decomposition of Calcium Carbonate (Limestone) (Equation 1.20)

• Reaction: $$CaCO_3(s) \xrightarrow{\text{Heat}} CaO(s) + CO_2(g)$$
  • Calcium carbonate (CaCO₃) decomposes to form calcium oxide (CaO) (quick lime) and carbon dioxide (CO₂) gas.
  • Application:
    • Quick lime (CaO) is used in the manufacture of cement.

Example 3: Decomposition of Lead Nitrate (Activity 1.6)

• Experiment:
  • Take lead nitrate (Pb(NO₃)₂) in a boiling tube.
  • Heat it over a flame.
  • Observation:
    • Brown fumes are released.
  • Reaction: $$2Pb(NO_3)_2(s) \xrightarrow{\text{Heat}} 2PbO(s) + 4NO_2(g) + O_2(g)$$
    • Lead nitrate (Pb(NO₃)₂) decomposes into lead oxide (PbO) (a yellow solid), nitrogen dioxide (NO₂) (brown fumes), and oxygen (O₂) gas.
    • Application:
      • This reaction is used to produce oxygen gas in laboratories.

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2. Photodecomposition (Decomposition by Light - Photolysis)

• Definition: When a compound breaks down in the presence of light, it is called photodecomposition or photolysis.
• This principle is used in black and white photography.

Example 1: Decomposition of Silver Chloride (Activity 1.7)

• Experiment:
  • Take silver chloride (AgCl) in a china dish.
  • Place it in sunlight for some time.
  • Observation:
    • White silver chloride turns grey.
  • Reaction: $$2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_2(g)$$
    • Silver chloride (AgCl) breaks down into silver metal (Ag) (grey in color) and chlorine (Cl₂) gas.

Example 2: Decomposition of Silver Bromide

• Reaction: $$2AgBr(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Br_2(g)$$
  • Silver bromide (AgBr) decomposes in sunlight to give silver metal and bromine gas.
  • Application:
    • This reaction is used in black and white photography.

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3. Electrolytic Decomposition (Decomposition by Electricity - Electrolysis)

• Definition: When a compound breaks down into simpler substances using electricity, it is called electrolytic decomposition or electrolysis.

Example: Electrolysis of Water (Activity 1.8)

• Experiment:
  • Take a plastic mug and drill two holes at its base.
  • Insert carbon electrodes into these holes.
  • Connect the electrodes to a 6V battery.
  • Fill the mug with water, adding a few drops of dilute sulphuric acid.
  • Invert two test tubes filled with water over the electrodes.
  • Switch on the current.
  • Observation:
    • Bubbles appear at both electrodes.
    • Water is displaced from the test tubes.
    • More gas is collected at one electrode than the other.
  • Reaction: $$2H_2O(l) \xrightarrow{\text{Electricity}} 2H_2(g) + O_2(g)$$
    • Water (H₂O) decomposes into hydrogen (H₂) and oxygen (O₂) gases.
    • More hydrogen gas is collected because water contains twice as many hydrogen atoms as oxygen atoms.
  • Testing the Gases:
    • Hydrogen test: A burning candle produces a "pop" sound when brought near the test tube containing hydrogen gas.
    • Oxygen test: A glowing splint relights in oxygen gas.

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Key Takeaways:

1. Decomposition Reactions break down one reactant into two or more simpler products.
2. They require energy input in the form of heat, light, or electricity and are endothermic.
3. Types of Decomposition Reactions:
   • Thermal Decomposition: Requires heat (e.g., decomposition of calcium carbonate).
   • Photodecomposition: Requires light (e.g., decomposition of silver chloride).
   • Electrolytic Decomposition: Requires electricity (e.g., electrolysis of water).
4. Applications:
   • Production of quick lime (CaO) for cement.
   • Black and white photography (AgCl and AgBr decomposition).
   • Oxygen production in labs (Pb(NO₃)₂ decomposition).
   • Electrolysis of water to obtain hydrogen and oxygen gases.

Questions And Answer 

1. A solution of a substance ‘X’ is used for whitewashing.

(i) Name the substance ‘X’ and write its formula.

• The substance ‘X’ used for whitewashing is calcium oxide (also called quick lime).
• Formula: CaO

(ii) Write the reaction of the substance ‘X’ named in (i) above with water.

When calcium oxide (CaO) reacts with water (H₂O), it forms calcium hydroxide (slaked lime, Ca(OH)₂) and releases heat.

$$CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + \text{Heat}$$

• Calcium hydroxide solution is used for whitewashing.
• It reacts slowly with carbon dioxide (CO₂) in the air to form a thin layer of calcium carbonate (CaCO₃), which gives the walls a shiny finish.

$$Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$$

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2. Why is the amount of gas collected in one of the test tubes in Activity 1.8 double of the amount collected in the other? Name this gas.

• Activity 1.8 refers to the electrolysis of water (H₂O), where water decomposes into hydrogen (H₂) and oxygen (O₂) gases when electricity is passed through it.

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• Reason: Water (H₂O) contains two hydrogen atoms for every one oxygen atom.

  • Hydrogen (H₂) gas is produced in double the amount compared to oxygen (O₂) gas.
  • This is why the gas collected in one test tube (hydrogen) is twice the amount of the gas collected in the other (oxygen).

• Name of the Gas: The gas collected in double amount is Hydrogen (H₂).

• Test for Hydrogen Gas: When a burning candle is brought near the test tube containing hydrogen, it burns with a ‘pop’ sound.

1.2.3 Displacement Reaction

Definition:

A displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its compound.

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Activity 1.9: Iron Nails in Copper Sulphate Solution

Procedure:

1. Take three iron nails and clean them using sandpaper.
2. Take two test tubes (A and B), and fill each with 10 mL of copper sulphate (CuSO₄) solution.
3. Tie two iron nails with a thread and immerse them in test tube B for about 20 minutes.
4. Keep one iron nail aside for comparison.
5. After 20 minutes, remove the nails from test tube B and observe:
   • Compare the blue colour intensity of copper sulphate solutions in test tubes A and B.
   • Compare the colour of iron nails dipped in the copper sulphate solution with the one kept aside.

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Observations:

• The iron nail turns brownish in colour.
• The blue colour of the copper sulphate solution fades.

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Explanation:

The following chemical reaction takes place:

$$Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$$

• Iron (Fe) is more reactive than copper (Cu).
• Iron displaces copper from copper sulphate (CuSO₄) solution, forming iron sulphate (FeSO₄).
• The blue colour of CuSO₄ solution fades because Cu²⁺ ions are replaced by Fe²⁺ ions, which form pale green iron sulphate solution.
• The iron nail turns brownish due to the deposition of copper (Cu) metal on its surface.

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Other Examples of Displacement Reactions:

1. Zinc (Zn) displaces copper from copper sulphate solution:

   $$Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$$
   • Zinc (Zn) is more reactive than copper (Cu).
   • Zinc replaces copper from CuSO₄, forming zinc sulphate (ZnSO₄).

2. Lead (Pb) displaces copper from copper chloride solution:

   $$Pb(s) + CuCl_2(aq) \rightarrow PbCl_2(aq) + Cu(s)$$
   • Lead (Pb) is more reactive than copper (Cu).
   • Lead replaces copper from CuCl₂, forming lead chloride (PbCl₂).

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Reactivity Series and Displacement Reactions:

• A more reactive metal can displace a less reactive metal from its salt solution.
• Reactivity order (from most to least reactive):
  K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag > Au
• Since iron (Fe), zinc (Zn), and lead (Pb) are more reactive than copper (Cu), they can displace Cu from its compounds.

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Conclusion:

• Displacement reactions occur when a more reactive element replaces a less reactive element in a compound.
• These reactions follow the reactivity series of metals.
• Iron, zinc, and lead can displace copper from its solutions, forming new compounds.
• The colour change in solutions indicates the formation of a new substance.

1.2.4 Double Displacement Reaction

Definition:

A double displacement reaction is a chemical reaction in which the ions of two compounds exchange places to form two new compounds. One of the products is often an insoluble precipitate, a gas, or water.

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Activity 1.10: Reaction Between Sodium Sulphate and Barium Chloride

Procedure:

1. Take 3 mL of sodium sulphate (Na₂SO₄) solution in a test tube.
2. Take 3 mL of barium chloride (BaCl₂) solution in another test tube.
3. Mix the two solutions together.
4. Observe the changes in the solution.

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Observation:

• A white, insoluble substance (precipitate) is formed in the solution.
• The formation of a precipitate indicates a chemical reaction has taken place.

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Explanation:

The reaction that takes place is:

$$Na_2SO_4 (aq) + BaCl_2 (aq) \rightarrow BaSO_4 (s) + 2NaCl (aq)$$

• Sodium sulphate (Na₂SO₄) and barium chloride (BaCl₂) react with each other.
• Barium sulphate (BaSO₄) is formed as a white precipitate, which is insoluble in water.
• Sodium chloride (NaCl) remains soluble in water.

The Ba²⁺ ions from BaCl₂ react with SO₄²⁻ ions from Na₂SO₄ to form BaSO₄, which precipitates out of the solution.

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Why is this a Double Displacement Reaction?

• In this reaction, the ions in the reactants are exchanged:
  • Ba²⁺ (from BaCl₂) replaces Na⁺ (from Na₂SO₄)
  • SO₄²⁻ (from Na₂SO₄) replaces Cl⁻ (from BaCl₂)
• As a result, two new compounds are formed:
  • Barium sulphate (BaSO₄) – precipitate
  • Sodium chloride (NaCl) – remains in solution

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Characteristics of Double Displacement Reactions:

1. Exchange of ions occurs between the reactants.
2. A precipitate, gas, or water is often formed.
3. These reactions commonly occur in aqueous solutions.
4. These reactions are also called precipitation reactions if a precipitate is formed.

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Other Examples of Double Displacement Reactions:

1. Reaction between silver nitrate and sodium chloride:

   $$AgNO_3 (aq) + NaCl (aq) \rightarrow AgCl (s) + NaNO_3 (aq)$$
   • Silver chloride (AgCl) forms a white precipitate.
   • Sodium nitrate (NaNO₃) remains soluble in water.

2. Reaction between lead nitrate and potassium iodide:

   $$Pb(NO_3)_2 (aq) + 2KI (aq) \rightarrow PbI_2 (s) + 2KNO_3 (aq)$$
   • Lead iodide (PbI₂) forms a yellow precipitate.
   • Potassium nitrate (KNO₃) remains in solution.

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Conclusion:

• Double displacement reactions involve the exchange of ions between two reactants.
• These reactions often produce a precipitate, gas, or neutral solution (water).
• The reaction between sodium sulphate and barium chloride produces an insoluble barium sulphate precipitate, making it a precipitation reaction.

1.2.5 Oxidation and Reduction (Redox Reactions)

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Activity 1.11: Oxidation of Copper Powder

Procedure:

1. Take about 1 g of copper powder in a china dish.
2. Heat it over a flame.
3. Observe the change in color of copper powder.

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Observation:

• The brown copper powder turns black.

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Explanation:

• The black substance formed is copper(II) oxide (CuO).

• This happens because oxygen is added to copper, forming copper oxide.

• The chemical reaction is:

  $$2Cu + O_2 \xrightarrow{\text{Heat}} 2CuO$$

• Here, copper (Cu) gains oxygen and is oxidized to copper(II) oxide (CuO).

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Reduction of Copper(II) Oxide

• If hydrogen gas (H₂) is passed over heated copper(II) oxide (CuO), the black CuO turns brown again, indicating that copper is obtained back.

• The chemical reaction is:

  $$CuO + H_2 \xrightarrow{\text{Heat}} Cu + H_2O$$

• Here:

  • CuO loses oxygen, so it is reduced to Cu.
  • H₂ gains oxygen, so it is oxidized to H₂O.

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Definition of Oxidation and Reduction

• Oxidation: Gaining oxygen or losing hydrogen.
• Reduction: Losing oxygen or gaining hydrogen.

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Redox Reactions (Oxidation-Reduction Reactions)

• A reaction in which one substance gets oxidized while the other gets reduced is called a Redox reaction.

Examples of Redox Reactions:

1. Zinc Oxide and Carbon Reaction:

   $$ZnO + C \rightarrow Zn + CO$$
   • ZnO loses oxygen (Reduced to Zn).
   • C gains oxygen (Oxidized to CO).

2. Manganese Dioxide and Hydrochloric Acid Reaction:

   $$MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$$
   • MnO₂ loses oxygen (Reduced to MnCl₂).
   • HCl gains oxygen (Oxidized to Cl₂).

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Conclusion:

• Oxidation and Reduction always occur together in a reaction.
• One substance gains oxygen (oxidation), while the other loses oxygen (reduction).
• These reactions are important in many industrial and biological processes.

1.3 EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE

Oxidation reactions are not just limited to laboratories; they also have significant effects in our daily lives. Two common examples of oxidation reactions affecting us are corrosion and rancidity.

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1.3.1 Corrosion

What is Corrosion?

• Corrosion is the gradual destruction of metals due to their reaction with moisture, air, acids, or other environmental factors.
• When metals react with oxygen and water in the environment, they form undesirable compounds that weaken them over time.

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Examples of Corrosion:

1. Iron (Fe)

   • Appearance: Reddish-brown rust
   • Chemical Reaction: $$4Fe + 3O_2 + 6H_2O \rightarrow 4Fe(OH)_3$$
   • The formation of iron hydroxide leads to the reddish-brown rust seen on iron objects.

2. Silver (Ag)

   • Appearance: Black coating (Silver sulfide)
   • Chemical Reaction: $$2Ag + H_2S \rightarrow Ag_2S + H_2$$
   • Silver reacts with hydrogen sulfide in the air, forming black silver sulfide, which tarnishes silver objects.

3. Copper (Cu)

   • Appearance: Green coating (Copper carbonate)
   • Chemical Reaction: $$2Cu + O_2 + CO_2 + H_2O \rightarrow Cu(OH)_2 \cdot CuCO_3$$
   • Copper reacts with oxygen, carbon dioxide, and water to form a green patina of copper carbonate, commonly seen on old copper roofs and statues.

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Effects of Corrosion:

• Weakens metal structures (e.g., bridges, buildings, vehicles).
• Reduces the lifespan of iron-based materials.
• Causes financial losses due to maintenance and replacement.

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How to Prevent Corrosion?

1. Painting, Greasing, or Oiling: Creates a protective barrier against air and moisture.
2. Galvanization: Coating iron with a layer of zinc to prevent rusting.
3. Alloying: Mixing metals (e.g., stainless steel) makes them more resistant to corrosion.
4. Electroplating: A thin layer of a less reactive metal is applied to prevent direct exposure to air.

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1.3.2 Rancidity

What is Rancidity?

• Rancidity occurs when fats and oils in food react with oxygen and undergo oxidation, leading to an unpleasant smell and taste.

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Causes of Rancidity:

1. Exposure to Oxygen: Leads to the breakdown of fats and oils.
2. High Temperature: Speeds up oxidation.
3. Presence of Light: Triggers chemical reactions that spoil the food.

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How to Prevent Rancidity?

1. Adding Antioxidants: Prevent oxidation of fats (e.g., Vitamin C, Vitamin E).
2. Airtight Storage: Keeps food away from oxygen.
3. Refrigeration: Slows down oxidation reactions.
4. Using Nitrogen Gas: Chips manufacturers flush packets with nitrogen gas to displace oxygen and prevent oxidation.

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Conclusion:

• Oxidation reactions affect our daily life in both negative (corrosion, rancidity) and positive (combustion, metabolism) ways.
• Understanding and preventing oxidation-related problems can save resources and enhance the durability of materials.

Questions and Answer 

1. Why does the colour of copper sulphate solution change when an iron nail is dipped in it?

   • When an iron nail is dipped in copper sulphate (CuSO₄) solution, a displacement reaction occurs. Iron, being more reactive than copper, displaces copper from the solution. As a result, the blue colour of copper sulphate fades, and a brownish layer of copper gets deposited on the iron nail.
   • Reaction: $$Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$$
   • The solution turns light green due to the formation of iron(II) sulphate (FeSO₄).

2. Give an example of a double displacement reaction other than the one given in Activity 1.10.

   • Example: Reaction between silver nitrate (AgNO₃) and sodium chloride (NaCl).
   • Reaction: $$AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$$
   • In this reaction, silver chloride (AgCl) forms as a white precipitate, and sodium nitrate (NaNO₃) remains in the solution.

3. Identify the substances that are oxidised and the substances that are reduced in the following reactions:

   (i) $4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$

   • Sodium (Na) is oxidised because it gains oxygen to form sodium oxide.
   • Oxygen (O₂) is reduced because it combines with sodium, meaning it loses its free molecular form.

   (ii) $CuO(s) + H_2(g) \rightarrow Cu(s) + H_2O(l)$

   • Copper(II) oxide (CuO) is reduced to copper (Cu) because it loses oxygen.
   • Hydrogen (H₂) is oxidised to water (H₂O) because it gains oxygen.

Answers:

1. Which of the statements about the reaction below are incorrect?

$$2PbO(s) + C(s) \rightarrow 2Pb(s) + CO_2(g)$$

(a) Lead is getting reduced. ✅ (Correct)
(b) Carbon dioxide is getting oxidised. ❌ (Incorrect) (CO₂ is not further oxidised)
(c) Carbon is getting oxidised. ✅ (Correct)
(d) Lead oxide is getting reduced. ✅ (Correct)

Correct answer: (i) (a) and (b)

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2. Identify the type of reaction in the given equation.

$$Fe_2O_3 + 2Al \rightarrow Al_2O_3 + 2Fe$$

This is a displacement reaction because aluminium displaces iron from iron(III) oxide.

Correct answer: (d) displacement reaction.

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3. What happens when dilute hydrochloric acid is added to iron filings?

$$Fe + 2HCl \rightarrow FeCl_2 + H_2$$

Correct answer: (a) Hydrogen gas and iron chloride are produced.

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4. What is a balanced chemical equation? Why should chemical equations be balanced?

A balanced chemical equation has an equal number of atoms of each element on both reactant and product sides. Balancing ensures the law of conservation of mass is followed.

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5. Translate the following statements into chemical equations and balance them.

(a) Hydrogen gas combines with nitrogen to form ammonia.

$$N_2 + 3H_2 \rightarrow 2NH_3$$

(b) Hydrogen sulphide gas burns in air to give water and sulfur dioxide.

$$2H_2S + 3O_2 \rightarrow 2H_2O + 2SO_2$$

(c) Barium chloride reacts with aluminium sulphate to give aluminium chloride and a precipitate of barium sulphate.

$$3BaCl_2 + Al_2(SO_4)_3 \rightarrow 2AlCl_3 + 3BaSO_4$$

(d) Potassium metal reacts with water to give potassium hydroxide and hydrogen gas.

$$2K + 2H_2O \rightarrow 2KOH + H_2$$

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6. Balance the following chemical equations.

(a) Nitric acid + Calcium hydroxide → Calcium nitrate + Water

$$2HNO_3 + Ca(OH)_2 \rightarrow Ca(NO_3)_2 + 2H_2O$$

(b) Sodium hydroxide + Sulfuric acid → Sodium sulfate + Water

$$2NaOH + H_2SO_4 \rightarrow Na_2SO_4 + 2H_2O$$

(c) Sodium chloride + Silver nitrate → Silver chloride + Sodium nitrate

$$NaCl + AgNO_3 \rightarrow AgCl + NaNO_3$$

(d) Barium chloride + Sulfuric acid → Barium sulfate + Hydrogen chloride

$$BaCl_2 + H_2SO_4 \rightarrow BaSO_4 + 2HCl$$

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7. Write the balanced chemical equations for the following reactions.

(a) Calcium hydroxide + Carbon dioxide → Calcium carbonate + Water

$$Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O$$

(b) Zinc + Silver nitrate → Zinc nitrate + Silver

$$Zn + 2AgNO_3 \rightarrow Zn(NO_3)_2 + 2Ag$$

(c) Aluminium + Copper chloride → Aluminium chloride + Copper

$$2Al + 3CuCl_2 \rightarrow 2AlCl_3 + 3Cu$$

(d) Barium chloride + Potassium sulphate → Barium sulphate + Potassium chloride

$$BaCl_2 + K_2SO_4 \rightarrow BaSO_4 + 2KCl$$

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8. Write the balanced chemical equation and identify the type of reaction.

(a) Potassium bromide + Barium iodide → Potassium iodide + Barium bromide

$$2KBr + BaI_2 \rightarrow 2KI + BaBr_2$$

Double displacement reaction

(b) Zinc carbonate → Zinc oxide + Carbon dioxide

$$ZnCO_3 \rightarrow ZnO + CO_2$$

Decomposition reaction

(c) Hydrogen + Chlorine → Hydrogen chloride

$$H_2 + Cl_2 \rightarrow 2HCl$$

Combination reaction

(d) Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen

$$Mg + 2HCl \rightarrow MgCl_2 + H_2$$

Displacement reaction

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9. Define exothermic and endothermic reactions with examples.

• Exothermic reaction: Releases heat.
  Example: Combustion of methane:

  $$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O + Heat$$

• Endothermic reaction: Absorbs heat.
  Example: Decomposition of calcium carbonate:

  $$CaCO_3 \xrightarrow{Heat} CaO + CO_2$$

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10. Why is respiration considered an exothermic reaction?

Respiration releases energy as glucose is broken down:

$$C_6H_{12}O_6 + 6O_2 \rightarrow 6CO_2 + 6H_2O + Energy$$

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11. Why are decomposition reactions opposite to combination reactions?

• Decomposition reaction: A compound breaks into simpler substances. $$2HgO \xrightarrow{Heat} 2Hg + O_2$$
• Combination reaction: Two or more substances combine. $$2H_2 + O_2 \rightarrow 2H_2O$$

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12. Examples of decomposition reactions using heat, light, or electricity.

13. Difference between displacement and double displacement reactions.

• Displacement: More reactive metal replaces a less reactive one. $$Zn + CuSO_4 \rightarrow ZnSO_4 + Cu$$
• Double displacement: Exchange of ions between compounds. $$BaCl_2 + Na_2SO_4 \rightarrow BaSO_4 + 2NaCl$$

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14. Recovery of silver from silver nitrate solution.

$$Cu + 2AgNO_3 \rightarrow Cu(NO_3)_2 + 2Ag$$

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15. What is a precipitation reaction?

A reaction that forms an insoluble solid.
Example:

$$BaCl_2 + Na_2SO_4 \rightarrow BaSO_4(s) + 2NaCl$$

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16. Define oxidation and reduction with examples.

(a) Oxidation: Gain of oxygen.

$$2Mg + O_2 \rightarrow 2MgO$$

(b) Reduction: Loss of oxygen.

$$CuO + H_2 \rightarrow Cu + H_2O$$

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17. Identify 'X' and the black compound.

• X = Copper (Cu)
• Compound = Copper(II) oxide (CuO)

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18. Why apply paint on iron articles?

To prevent rusting by blocking oxygen and moisture.

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19. Why are food items flushed with nitrogen?

To prevent oxidation and rancidity.

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20. Define corrosion and rancidity.

(a) Corrosion: Metal deterioration (e.g., rusting of iron).
(b) Rancidity: Spoilage of fats and oils due to oxidation.