Chemical Reactions and Equations Notes Class 10 Science Chapter 1

Class 10 Science • Chapter 1

Chemical Reactions and Equations

Complete Notes based on NCERT

Complete NCERT-based notes for Class 10 Science Chapter 1 covering chemical reactions, types, activities, and equations.

1. Chemical Reaction

A chemical reaction is a process in which one or more substances (reactants) change into new substances (products) with different properties.

Examples:
  • Milk → Curd
  • Rusting of iron
  • Fermentation of grapes
  • Cooking of food
  • Digestion

Types of Change in Matter

1. Physical Change

Only physical properties change. No new substance is formed.

Ice → Water → Vapour

2. Chemical Change

New substance is formed.

  • Rusting
  • Burning of paper

Activity 1.1 – Magnesium Ribbon

Procedure

  • Clean ribbon with sandpaper
  • Burn using tongs
  • Collect ash

Observation

  • Dazzling white flame
  • White powder (MgO)
2Mg + O₂ → 2MgO

Activity 1.2 – Lead Nitrate + Potassium Iodide

Procedure

  • Take lead nitrate solution in test tube
  • Add potassium iodide solution

Observation

  • Yellow precipitate (PbI₂)
Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃

Activity 1.3 – Zinc with Acid

Procedure

  • Take zinc granules in flask
  • Add dilute HCl or H₂SO₄
  • Touch flask carefully

Observation

  • Hydrogen gas bubbles
  • Flask becomes warm
  • Pop sound confirms hydrogen
Zn + 2HCl → ZnCl₂ + H₂ ↑
Zn + H₂SO₄ → ZnSO₄ + H₂

Characteristics of Chemical Reaction

  • Formation of precipitate
  • Evolution of gas
  • Colour change
  • Temperature change
  • State change

Chemical Equation

A chemical equation represents a reaction using symbols and formulas.

2H₂ + O₂ → 2H₂O

Balanced Chemical Equation

A balanced equation has equal number of atoms on both sides.

A₂ + B₂ → 2AB
Start Test

Why do we balance chemical equations?

Answer:

  • Because of the Law of Conservation of Mass
  • Mass can neither be created nor destroyed
  • Total mass of reactants is equal to total mass of products

Example:

3Fe + 4H₂O → Fe₃O₄ + 4H₂

Explanation:

Fe₃O₄ contains both Fe²⁺ and Fe³⁺ states (mixed oxide of iron).
CaCO₃ → CaO + CO₂

Numerical Problem

Consider the following chemical equation:

xAl + yH₂O → zAl₂O₃ + nH₂

Find x, y, z and n

2Al + 3H₂O → Al₂O₃ + 3H₂
x = 2,   y = 3,   z = 1,   n = 3

Balancing Chemical Equation Example

In order to balance:

MnO₂ + xHCl → MnCl₂ + yH₂O + zCl₂

Find x, y and z

MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
x = 4,   y = 2,   z = 1

Question

Why do we balance a chemical equation? Name and state the law that suggests the balancing of chemical equation and also balance the equation:

Zn + H₃PO₄ → Zn₃(PO₄)₂ + H₂

Answer:

A chemical equation is balanced to ensure that the number of atoms of each element is equal on both sides (reactants = products). This is important because atoms are neither created nor destroyed during a chemical reaction.

Law behind balancing equations

Balancing of chemical equations is based on the Law of Conservation of Mass.

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

So, the total mass (and number of atoms) of reactants must be equal to that of products.

Balancing the given equation

Unbalanced equation:

Zn + H₃PO₄ → Zn₃(PO₄)₂ + H₂

Step-by-step balancing:

Step 1: Balance Zn
Right side has 3 Zn → put 3Zn on left
3Zn + H₃PO₄ → Zn₃(PO₄)₂ + H₂
Step 2: Balance PO₄
Right side has 2 PO₄ → put 2H₃PO₄
3Zn + 2H₃PO₄ → Zn₃(PO₄)₂ + H₂
Step 3: Balance Hydrogen
Balance H atoms → put coefficient 3 before H₂
3Zn + 2H₃PO₄ → Zn₃(PO₄)₂ + 3H₂

Final Balanced Equation:

3Zn + 2H₃PO₄ → Zn₃(PO₄)₂ + 3H₂

Types of Chemical Reaction

  • Combination Reaction
  • Decomposition Reaction
  • Displacement Reaction
  • Double Displacement Reaction
  • Exothermic and Endothermic Reaction
  • Oxidation and Reduction

Types of Chemical Reactions (Detailed)

1. Combination Reaction

A reaction in which two or more reactants combine to form a single product is called a combination reaction.
A + B → AB

Example:

CaO + H₂O → Ca(OH)₂ + Heat
Quicklime (CaO) reacts with water to form slaked lime [Ca(OH)₂]. This reaction is highly exothermic (produces heat).
Ca(OH)₂ + CO₂ → CaCO₃ + H₂O
Calcium hydroxide is used for whitewashing walls. It reacts with carbon dioxide to form calcium carbonate, which gives a shiny finish.

More Examples:

C + O₂ → CO₂
2H₂ + O₂ → 2H₂O
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2. Decomposition Reaction

A reaction in which a single reactant breaks down to form two or more products is called a decomposition reaction. This reaction is opposite to combination reaction.

Types of Decomposition Reaction

A. Thermal Decomposition

This reaction uses heat energy for decomposition of the reactant.
example1- CaCO₃ → CaO + CO₂
Calcium oxide (CaO) is used in the manufacturing of cement.

Ex-2 Decomposition of Ferrous Sulphate

On heating ferrous sulphate crystals, they decompose to form ferric oxide, sulphur dioxide and sulphur trioxide.
FeSO₄·7H₂O → FeSO₄ + 7H₂O
2FeSO₄ → Fe₂O₃ + SO₂ + SO₃
Fe₂O₃ is solid, SO₂ and SO₃ are gases.

Ex3- Decomposition of Lead Nitrate

On heating lead nitrate, it decomposes to give lead oxide (yellow), nitrogen dioxide (brown fumes) and oxygen gas.
2Pb(NO₃)₂ → 2PbO + 4NO₂ + O₂

B . Electrolysis

This reaction involves the use of electrical energy for decomposition of reactant molecules.

Electrolysis of Water

When electric current is passed through water, it decomposes to give hydrogen and oxygen gas.
2H₂O → 2H₂ + O₂
At Anode (+): Oxygen gas (O₂) is formed (oxidation)
At Cathode (−): Hydrogen gas (H₂) is formed (reduction)
Ratio of gases produced:
Hydrogen : Oxygen = 2 : 1
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Electrolysis of Molten Sodium Chloride

When electric current is passed through molten sodium chloride (NaCl), it decomposes to give sodium metal and chlorine gas.
2NaCl → 2Na + Cl₂
At Anode: 2Cl⁻ → Cl₂
At Cathode: Na⁺ → Na
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C . Photolysis (Photochemical Decomposition)

This reaction involves the use of light energy for decomposition.

Example 1:

2AgCl → 2Ag + Cl₂
Silver chloride decomposes in sunlight to form silver (greyish white) and chlorine gas (greenish yellow).

Example 2:

2AgBr → 2Ag + Br₂
This reaction is used in black and white photography.

3. Displacement Reaction

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

Reactivity Series

K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Hg > Ag > Au > Pt

Types of Displacement

(a) Single Displacement Reaction

It is a type of chemical reaction where an element reacts with a compound and takes the place of another element in that compound.

Example:

Zn + CuSO₄ → ZnSO₄ + Cu
Zinc is more reactive than copper, so it displaces copper from copper sulphate solution.
Cu + ZnSO₄ → No Reaction
Copper is less reactive than zinc, so it cannot displace zinc from zinc sulphate.

More Examples of Displacement Reaction

Zinc is more reactive than copper, so it displaces copper from CuSO₄ solution to form zinc sulphate and copper metal.
Pb + CuCl₂ → PbCl₂ + Cu
Fe + CuSO₄ → FeSO₄ + Cu
Blue colour of CuSO₄ solution changes to green due to formation of FeSO₄.
Fe + H₂O → Fe₃O₄ + H₂
AgNO₃ + Cu → Cu(NO₃)₂ + Ag
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4. Double Displacement Reaction

A reaction in which two different ions or groups of atoms in the reactant molecules are exchanged to form new products is called a double displacement reaction.
Sometimes, it is also called a precipitation reaction because a precipitate is formed.

Example:

On adding sodium sulphate solution to barium chloride solution, a white precipitate of barium sulphate is formed.
Na₂SO₄ + BaCl₂ → 2NaCl + BaSO₄ ↓
BaSO₄ is a white precipitate.

More Examples of Double Displacement Reaction

On adding silver nitrate solution to sodium bromide solution, a yellow precipitate of silver bromide is formed along with sodium nitrate.
AgNO₃ + NaBr → AgBr ↓ + NaNO₃
AgBr is a yellow precipitate.
KI + Pb(NO₃)₂ → 2KNO₃ + PbI₂ ↓
PbI₂ is a yellow precipitate.
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5. Exothermic Reaction

Reactions in which heat is released are called exothermic reactions.

Examples:

1. Respiration is an exothermic process.
2. Decomposition of vegetable matter is exothermic.
3. Most combination reactions and combustion reactions are exothermic.

Burning of Natural Gas:

CH₄ + O₂ → CO₂ + H₂O + Heat
CₓHᵧ + (x + y/4)O₂ → xCO₂ + (y/2)H₂O + Heat
CH₄ + 2O₂ → CO₂ + 2H₂O + Heat

More Examples of Combustion (Exothermic Reaction)

2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O + Heat
C₂H₄ + 3O₂ → 2CO₂ + 2H₂O
2Mg + O₂ → 2MgO + Heat
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6. Endothermic Reaction

Reactions which occur by absorption of heat, light, or electrical energy are called endothermic reactions.

Examples:

1. Photosynthesis is an endothermic process.
2. Most decomposition reactions are endothermic.
(Note: Decomposition of vegetable matter into compost is exothermic.)
6CO₂ + 12H₂O → C₆H₁₂O₆ + 6O₂ + 6H₂O
(Occurs in presence of sunlight and chlorophyll)

More Examples:

NH₄Cl → NH₃ + HCl
HgO → Hg + O₂

Neutralisation Reaction

A reaction in which an acid reacts with a base to form salt and water is called a neutralisation reaction.
HCl + NaOH → NaCl + H₂O
Acid + Base → Salt + Water
H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O
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Oxidation

Oxidation is defined as:
  • Addition of oxygen to a substance
  • Removal of hydrogen from a substance
  • Loss of electrons
Mg + O₂ → MgO
Magnesium is oxidised to magnesium oxide.

Oxidation (Detailed)

Oxidation is defined as:
  • Addition of oxygen
  • Removal of hydrogen
  • Loss of electrons
Cu + O₂ → CuO
Oxygen is added to copper, so oxidation occurs.
H₂S + O₂ → S + H₂O
Hydrogen is removed, so oxidation occurs.
Zn → Zn²⁺ + 2e⁻
Loss of electrons means oxidation.
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Reduction

Reduction is defined as:
  • Removal of oxygen
  • Addition of hydrogen
  • Gain of electrons
2KClO₃ → 2KCl + 3O₂
Oxygen is removed from KClO₃, so reduction occurs.
2Na + H₂ → 2NaH
Hydrogen is added, so reduction occurs.
Zn²⁺ + 2e⁻ → Zn
Gain of electrons means reduction.

Important Notes

  • If a substance gains oxygen during a reaction, it is said to be oxidised.
  • If a substance loses oxygen during a reaction, it is said to be reduced.

Oxidising Agent

The substance which causes oxidation of another substance is called an oxidising agent.
(It itself gets reduced.)

Reducing Agent

The substance which causes reduction of another substance is called a reducing agent.
(It itself gets oxidised.)
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Redox Reaction

A reaction in which oxidation and reduction take place simultaneously is called a redox reaction.

Example 1

CuO + H₂ → Cu + H₂O
CuO is reduced to Cu.
H₂ is oxidised to H₂O.
Oxidising Agent: CuO
Reducing Agent: H₂
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Example 2

MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O
This reaction involves both oxidation and reduction processes.

Additional Redox Examples

MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O
MnO₂ is reduced to MnCl₂ (removal of oxygen).
HCl is oxidised to Cl₂ (removal of hydrogen).
Oxidising Agent: MnO₂
Reducing Agent: HCl
2H₂S + O₂ → 2S + 2H₂O
H₂S is oxidised (removal of hydrogen).
O₂ acts as oxidising agent.
CuO + C → Cu + CO
CuO is reduced (removal of oxygen).
C is oxidised.
Oxidising Agent: CuO
Reducing Agent: C
FeO + CO → Fe + CO₂
FeO is reduced to Fe.
CO is oxidised to CO₂.
Oxidising Agent: FeO
Reducing Agent: CO
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Effects of Oxidation in Everyday Life

Corrosion

Corrosion is the degradation of a metal due to the action of air, moisture, or chemicals on its surface.
Corrosion is an oxidation process.
Example: Iron articles get coated with reddish-brown powder when left in moist air for a long time. This process is called rusting of iron.

Rust and Rusting

Rust is hydrated ferric oxide: Fe₂O₃·xH₂O
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Prevention of Rusting

  • By painting
  • By greasing or oiling
  • By galvanisation
  • By tin plating and chromium plating
  • By alloying
Galvanisation is the process of depositing a thin layer of zinc on iron. Zinc reacts with oxygen to form zinc oxide, which forms a protective layer over iron and prevents corrosion.
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Corrosion of Silver and Copper

Silver and copper lose their lustre due to formation of:
  • Silver sulphide (black coating)
  • Basic copper carbonate (green coating)
2Ag + H₂S → Ag₂S + H₂
Silver reacts with hydrogen sulphide to form silver sulphide (black coating).
2Cu + CO₂ + O₂ + H₂O → CuCO₃·Cu(OH)₂
Copper reacts with air, moisture and carbon dioxide to form basic copper carbonate (green coating).

Rancidity

Rancidity is the process of slow oxidation of oils and fats present in food materials, which results in a change in smell and taste.

Methods to Prevent Rancidity

  • Keeping food in airtight containers
  • Storing food at low temperature (refrigeration)
  • Packing food in nitrogen gas (to prevent oxidation)
  • Avoid keeping cooked food in direct sunlight
  • Adding antioxidants to food