CBSE Class 10 Science Chemical Reactions and Equations Quick Revision for Board Exams

With the upcoming CBSE Class 10 Science Board Exam on February 25, 2026, students must focus on quick and effective revision of important chapters. Chemical Reactions and Equations is one of the most scoring and fundamental chapters in the syllabus. 

It forms the base for understanding advanced chemistry topics in higher classes and frequently carries direct theory and numericals in the board paper. Here is a revision guide to help students revise key concepts, important definitions, and commonly asked questions efficiently before the board exam.

Chemical Reactions and Their Characteristics

A chemical change is known as a chemical reaction. Several observable indicators suggest a chemical reaction has occurred:

1. Change in Color: A visible change in the color of the substances.

2. Formation of a Precipitate: The formation of an insoluble solid during a chemical reaction. A substance must be formed as a result of a chemical reaction to be called a precipitate.

3. Evolution of a Gas: The release of a gas.

4. Change in Temperature: The temperature may increase (Exothermic reaction) or decrease (Endothermic reaction).

5. Change in Physical State: A change from solid to liquid, liquid to gas, etc. This factor alone is not sufficient; multiple factors must be analyzed.

Balancing Chemical Equations

A Balanced Chemical Equation ensures the number of atoms of each element is equal on both the reactant and product sides.

Why Balance?

Balancing satisfies the Law of Conservation of Mass (Antoine Lavoisier, 1789), which states that mass is neither created nor destroyed in a chemical reaction. Thus, the total mass and number of atoms of each element of reactants must equal that of products.

How to Balance:

Balancing is achieved by adding or changing stoichiometric coefficients (numbers placed in front of chemical formulas). Never change the subscripts in a chemical formula, as this alters the substance's chemical identity.

• Example: 2Mg + O₂ → 2MgO (Reactants: 2 Mg, 2 O; Products: 2 Mg, 2 O)

Types of Chemical Reactions

In chemistry, reactions are broadly classified into different types based on how substances interact and transform.

Exothermic and Endothermic Reactions

These reactions are classified based on heat exchange with the surroundings.

(For Exothermic, think Ex- as in Exit and -thermic as in heat. Heat exits the reaction.)

1. Combination (Synthesis) Reaction

A reaction where multiple reactants combine to form a single product.

• General Form: A + B → AB

Examples:

1. Element + Element → Compound: 2Mg(s) + O₂(g) → 2MgO(s) + Heat + Light (Magnesium burns with a dazzling white flame to form white magnesium oxide).

Applications:

• Whitewashing: Slaked lime Ca(OH)₂ solution is used for whitewashing. It reacts with atmospheric carbon dioxide over 2-3 days to form a thin, hard layer of calcium carbonate, giving a shiny finish.

• Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

• CaCO₃ is known as Limestone, Marble, or Chalk.

2. Decomposition Reaction

A reaction where a single reactant breaks down into multiple simpler products. This is the opposite of a combination reaction and is always endothermic, requiring energy input.

• General Form: AB → A + B

Types of Decomposition:

1. Thermal Decomposition (Thermolysis): Decomposition by Heat.

• FeSO₄(s) --(Heat)--> Fe₂O₃(s) + SO₂(g) + SO₃(g) (Ferrous sulfate decomposes to reddish-brown ferric oxide and sulfur gases with a burning matchstick smell).

• CaCO₃(s) --(Heat)--> CaO(s) + CO₂(g) (Limestone decomposes into quicklime and carbon dioxide; quicklime is used in the cement industry).

• 2Pb(NO₃)₂(s) --(Heat)--> 2PbO(s) + 4NO₂(g) + O₂(g) (Lead nitrate (white) decomposes into yellow lead oxide, brown fumes of nitrogen dioxide, and oxygen gas).

2. Electrolytic Decomposition (Electrolysis): Decomposition by Electrical Energy.

• 2H₂O(l) --(Electricity)--> 2H₂(g) + O₂(g) (Water decomposes into hydrogen (at cathode) and oxygen (at anode)). The volume of hydrogen is twice that of oxygen. An electrolyte is added to increase water's conductivity.

3. Photolytic Decomposition (Photolysis): Decomposition by Light Energy.

• 2AgCl(s) --(Sunlight)--> 2Ag(s) + Cl₂(g)

• 2AgBr(s) --(Sunlight)--> 2Ag(s) + Br₂(g)

• Silver chloride (white) and silver bromide (pale yellow) decompose into silver (greyish-white) and halogen gas upon light exposure. This is used in black and white photography.

3. Displacement Reaction

A reaction where a more reactive element displaces a less reactive element from its compound in a solution. These are generally exothermic.

• General Form: A + BC → AC + B (A is more reactive than B).

• Reactivity is determined by the reactivity series.

Examples:

1. Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s) (Iron, more reactive than copper, displaces it. Blue copper sulfate turns pale green, and a reddish-brown copper layer deposits on iron).

4. Double Displacement (Metathesis) Reaction

A reaction involving an exchange of ions (cations and anions) between two reactant compounds.

• General Form: AB + CD → AD + CB

Types of Double Displacement Reactions:

1. Precipitation Reaction: A double displacement reaction where one product is an insoluble solid (precipitate).

• Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq) (Sodium sulfate and Barium chloride form a white precipitate of Barium sulfate).

• 2KI(aq) + Pb(NO₃)₂(aq) → PbI₂(s)↓ + 2KNO₃(aq) (Potassium iodide and Lead nitrate form a yellow precipitate of Lead iodide).

5. Redox Reactions

Reactions in which oxidation and reduction occur simultaneously.

Example: Oxidation of Copper

• 2Cu(s) + O₂(g) --(Heat)--> 2CuO(s)

• Copper gains oxygen to form black copper(II) oxide. Copper is oxidized.

Oxidizing and Reducing Agents

• Oxidizing Agent (Oxidant): Causes oxidation in another substance by getting reduced itself.

• Reducing Agent (Reductant): Causes reduction in another substance by getting oxidized itself.

In CuO + H₂ → Cu + H₂O:

• CuO is the oxidizing agent (gets reduced).

• H₂ is the reducing agent (gets oxidized).

Key Rule: The substance that undergoes reduction is the oxidizing agent. The substance that undergoes oxidation is the reducing agent.

Effects of Oxidation in Daily Life

These are practical examples of redox reactions.

1. Corrosion

Corrosion is the gradual degradation of metals attacked by atmospheric substances (gases, moisture, acids).

• Rusting of Iron: Exposure to moist air forms a flaky, reddish-brown layer of hydrated iron(III) oxide (rust).

• Tarnishing:

• Silver turns black (silver sulfide).

• Copper forms a green layer (copper carbonate and hydroxide).

2. Rancidity

Rancidity is the aerial oxidation of fats and oils in food items, leading to unpleasant smell and taste, making food unfit.

Prevention of Rancidity:

1. Using Antioxidants: Adding substances that prevent oxidation (e.g., Vitamin C or E).

2. Flushing with Nitrogen Gas: Replacing oxygen in packaged foods with inert nitrogen.

3. Refrigeration: Low temperatures slow oxidation.

4. Airtight Containers: Limits exposure to atmospheric oxygen.