Notes Class 10 Science Chapter 1 Chemical Reactions and Equations

Notes For All Chapters Science Class 10 CBSE

1. Introduction

In daily life we see many changes:

• Milk turns sour in summer.
• Iron objects rust in moist air.
• Grapes ferment to form alcohol.
• Cooking and digestion of food.
• Respiration in our body.

In all these, the nature and identity of substances change.

These are chemical changes → involve chemical reactions.

2. How to Identify a Chemical Reaction

A chemical reaction has occurred if you observe:

• Change in state.
• Change in colour.
• Evolution of a gas.
• Change in temperature.

3. Chemical Equations

3.1 Word Equations

Example:

• Magnesium + Oxygen → Magnesium oxide
• Reactants → Substances that undergo change.
• Products → New substances formed.

3.2 Skeletal Equation

Using chemical formulae:

• Mg + O₂ → MgO

If number of atoms on LHS ≠ RHS → equation is unbalanced.

3.3 Balanced Chemical Equations

Law of Conservation of Mass → Mass can neither be created nor destroyed.

Therefore, atoms of each element must be equal on both sides.

Example:

• Zn + H₂SO₄ → ZnSO₄ + H₂ (Balanced)

Balancing steps use hit-and-trial method.

3.4 Physical States in Equations

(s) → solid, (l) → liquid, (g) → gas, (aq) → aqueous solution.

Example:

• 3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

3.5 Conditions of Reactions

Temperature, pressure, catalyst, etc. can be written above/below the arrow.

Example:

• CO + 2H₂ → CH₃OH (in presence of catalyst, pressure).

4. Types of Chemical Reactions

4.1 Combination Reaction

• Two or more substances combine to form one product.
• Example:
• CaO + H₂O → Ca(OH)₂ (slaked lime)
• C + O₂ → CO₂
• 2H₂ + O₂ → 2H₂O
• Often exothermic (release heat).
• Respiration is also exothermic:
• C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + energy

4.2 Decomposition Reaction

One substance breaks down to form two or more products.

Thermal decomposition (by heat):

• 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃
• CaCO₃ → CaO + CO₂

Electrolytic decomposition (by electricity):

• 2H₂O → 2H₂ + O₂

Photodecomposition (by light):

• 2AgCl → 2Ag + Cl₂ (in sunlight)
• Used in black & white photography.

Decomposition reactions often need energy → endothermic.

4.3 Displacement Reaction

A more reactive element displaces a less reactive element from its compound.

Examples:

• Fe + CuSO₄ → FeSO₄ + Cu
• Zn + CuSO₄ → ZnSO₄ + Cu
• Pb + CuCl₂ → PbCl₂ + Cu

4.4 Double Displacement Reaction

Exchange of ions between two compounds.

Example (precipitation reaction):

• Na₂SO₄ + BaCl₂ → BaSO₄ (white ppt) + 2NaCl

4.5 Oxidation and Reduction (Redox Reactions)

Oxidation → Gain of oxygen / Loss of hydrogen.

Reduction → Loss of oxygen / Gain of hydrogen.

Example:

• 2Cu + O₂ → 2CuO (Oxidation of Cu)
• CuO + H₂ → Cu + H₂O (Reduction of CuO, Oxidation of H₂)

In redox reactions, one substance gets oxidised, the other gets reduced.

5. Effects of Oxidation Reactions in Daily Life

5.1 Corrosion

Metals react with substances around (air, water, acids).

Examples:

• Rusting of iron (reddish brown coating).
• Black coating on silver.
• Green coating on copper.

Corrosion causes damage to bridges, ships, railings, etc.

5.2 Rancidity

Oxidation of fats and oils makes food smell/taste bad.

Prevented by:

• Adding antioxidants.
• Storing in airtight containers.
• Flushing packets with nitrogen gas (e.g., chips packets).