Class 10 Chemical Reactions and Equations 2026 Board Exam Mind Map & Revision

Chemical Reactions and Equations: The chapter Chemical Reactions and Equations serves as the foundation of class 10th Chemistry section and is frequently tested through balancing tasks, observation-based questions, and redox analysis. With the Class 10 Science board exam scheduled for 25 February 2026, students are now in the most crucial phase of their revision. 

To make preparation easier and effective, this mind map–based explanation helps you quickly connect essential concepts. It will help candidates identify chemical changes and balancing equations and understand reaction types and redox processes in a clear, structured way. This explanation is designed for fast retention and high-scoring performance, ensuring you walk into your exam centre with confidence.

Class 10 Chemical Reactions and Equations  Full Chapter Mind Map Series

The chapter Chemical Reactions and Equations is the foundation of Class 10 Chemistry Syllabus and carries significant weight in board examinations. The concepts are logical, and students often find it challenging to master balancing equations, identifying types of redox reactions, and remembering specific colour changes. To simplify these transformations and make your exam revision highly effective, we have organised this content into a clear, well-structured flow.

Watch Video for Easy Explanation: Class 10 Chemical Reactions and Equations Mind Map

This comprehensive mind map video synthesizes the entire chapter into a clear, connected format, empowering students to grasp complex chemical concepts and revise with total confidence for their exams.

The Concept of Change

A change is defined as a process where the final state of a substance differs from its initial state. While various types of changes exist, a primary classification determines whether the chemical composition of the substance is altered.

Physical vs. Chemical Changes

Changes are categorized into two main types based on their effect on a substance's chemical makeup.

The most reliable way to differentiate between the two is to assess whether the chemical composition has been altered.

Chemical Reactions and Their Characteristics

The process through which a chemical change occurs is called a chemical reaction. Several visible indicators, known as the characteristics of a chemical reaction, help determine if a reaction has taken place.

These characteristics are:

1. Change in Color: The color of the substances may change.

2. Formation of a Precipitate: The formation of an insoluble solid (precipitate) when solutions are mixed.

3. Evolution of a Gas: A gas may be produced.

4. Change in Physical State: Reactants in one state may produce products in another.

• Caution: If a change in physical state is the only observation, it may indicate a physical change rather than a chemical one.

5. Change in Temperature: The temperature of the immediate surroundings of the reaction may increase or decrease.

Example of Precipitate Formation: When colorless solutions of Barium Chloride (BaCl₂) and Sodium Sulfate (Na₂SO₄) are mixed, a white precipitate of Barium Sulfate (BaSO₄) forms, along with a colorless solution of Sodium Chloride (NaCl).

Example of Gas Evolution & State Change: The combustion of charcoal (Carbon, a solid) in the presence of Oxygen (a gas) produces Carbon Dioxide (a gas), along with heat and light energy.

Change in Temperature: Exothermic vs. Endothermic Reactions

Chemical reactions often involve energy changes, typically observed as temperature changes.

• Exothermic Reactions: Energy is released into the surroundings, primarily as heat and/or light. (Memory Tip: Think "Exo" for Exit. Energy (heat) exits the system.) Examples include combustion, respiration, and composting.

• Endothermic Reactions: Energy is absorbed from the surroundings or must be supplied for the reaction. (Memory Tip: Think "Endo" for Entry. Energy (heat) enters the system.) Examples include photosynthesis (sunlight), thermolysis (heat), and electrolysis (electrical energy).

Representing Chemical Reactions

There are progressively more concise and informative ways to represent a chemical reaction.

1. Word Equation: Represents the reaction using substance names. Reactants are on the left, Products on the right.

• Example: Magnesium + Oxygen → Magnesium Oxide

2. Skeletal Chemical Equation: Uses chemical symbols and formulas. This is often an unbalanced chemical equation.

• Example: Mg + O₂ → MgO

3. Balanced Chemical Equation: An equation where the number of atoms of each element is equal on both reactant and product sides.

• Example: 2Mg + O₂ → 2MgO

Balancing Equations and Making Them Informative

The Need to Balance a Chemical Equation

Balancing chemical equations is essential to satisfy the Law of Conservation of Mass, which states that mass is neither created nor destroyed in a chemical reaction. Since atoms are only rearranged, the number of atoms of each element must remain constant before and after the reaction.

Making Chemical Equations More Informative

To provide more detail, chemical equations are enhanced with additional information:

1. Physical States:

• (s): Solid; (l): Liquid; (g): Gas; (aq): Aqueous solution (dissolved in water).

• ↑: Indicates an evolved gas.

• ↓ or (ppt): Indicates a precipitate (insoluble solid).

2. Reaction Conditions: Conditions like temperature, pressure, or catalysts are written above or below the arrow.

• Example (Pressure): CO(g) + 2H₂(g) --[340 atm]--> CH₃OH(l)

• Example (Light & Catalyst): 6CO₂(aq) + 6H₂O(l) --[Sunlight / Chlorophyll]--> C₆H₁₂O₆(aq) + 6O₂(g)

Types of Chemical Reactions

1. Combination (or Synthesis) Reaction

A reaction where two or more reactants (elements or compounds) combine to form a single product.

• General Form: A + B → AB

Example 1: Formation of Magnesium Oxide

• Reaction: 2Mg(s) + O₂(g) → 2MgO(s) + Heat + Light

• Context: Magnesium (Mg) is a reactive metal forming a protective magnesium oxide layer. Therefore, a magnesium ribbon must be cleaned with sandpaper before ignition.

• Observation: When heated, it burns with a dazzling white flame producing white magnesium oxide powder.

• Safety Precaution: This bright flame emits UV light, which can cause eye damage. Eye protection must be worn.

Example 2: Formation of Slaked Lime

• Reaction: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat

• Common Names: CaO is Quicklime; Ca(OH)₂ is Slaked Lime.

• Observations: This is a very fast and highly exothermic reaction. A hissing sound is produced as significant heat causes water to boil.

• Application in Whitewashing:

1. Slaked lime Ca(OH)₂ solution is applied to walls.

2. Over 2-3 days, it reacts slowly with carbon dioxide (CO₂) from the air to form a thin, hard, shiny layer of Calcium Carbonate (CaCO₃).
   Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

3. Common names for CaCO₃ include Limestone, Marble, and Chalk.

2. Decomposition Reaction

A reaction where a single reactant breaks down into two or more simpler products. It is the inverse of a combination reaction. Energy (heat, electricity, or light) must be supplied.

• General Form: AB → A + B

A. Thermal Decomposition (Thermolysis): Decomposition by Heat

• Example 1: Decomposition of Ferrous Sulfate

1. Heating pale green hydrated ferrous sulfate (FeSO₄·7H₂O) first forms white anhydrous ferrous sulfate (FeSO₄).

2. Further heating decomposes FeSO₄ into reddish-brown Ferric Oxide (Fe₂O₃) and two pungent gases: Sulfur Dioxide (SO₂) and Sulfur Trioxide (SO₃), smelling like a burning matchstick.
   2FeSO₄(s) --[Heat]--> Fe₂O₃(s) + SO₂(g) + SO₃(g)

• Safety Precautions: Do not point the test tube mouth towards anyone. Gently waft evolved gases to smell; do not sniff directly. Always use a test tube holder.

• Example 2: Decomposition of Calcium Carbonate

• Heating limestone (CaCO₃) produces quicklime (CaO) and carbon dioxide (CO₂). CaO is vital in cement manufacturing.
  CaCO₃(s) --[Heat]--> CaO(s) + CO₂(g)

• Example 3: Decomposition of Lead Nitrate

• Heating white lead nitrate (Pb(NO₃)₂) produces:

1. A yellow residue of Lead Oxide (PbO).

2. Brown fumes of Nitrogen Dioxide (NO₂) gas.

3. Colorless Oxygen (O₂) gas.

Observation: A crackling sound (decrepitation) is heard during heating.
2Pb(NO₃)₂(s) --[Heat]--> 2PbO(s) + 4NO₂(g) + O₂(g)

B. Electrolytic Decomposition (Electrolysis): Decomposition by Electricity

• Example: Electrolysis of Water

• Passing electric current through water decomposes it into hydrogen and oxygen gas.
  2H₂O(l) --[Electricity]--> 2H₂(g) + O₂(g)

• Observations & Key Points: Hydrogen gas (H₂) is collected at the cathode (negative electrode). Oxygen gas (O₂) is collected at the anode (positive electrode). The volume of hydrogen gas produced is double the volume of oxygen gas, reflecting the 2:1 ratio in water. Pure water is a poor conductor of electricity; an acid or salt is added as an electrolyte.

C. Photolytic Decomposition (Photolysis): Decomposition by Light

• Example: Decomposition of Silver Halides

• Silver Chloride (AgCl): This white solid decomposes into silvery-grey Silver (Ag) and greenish-yellow Chlorine (Cl₂) gas when exposed to sunlight.
  2AgCl(s) --[Sunlight]--> 2Ag(s) + Cl₂(g)

• Silver Bromide (AgBr): This yellow solid decomposes into silvery-grey Silver (Ag) and reddish-brown Bromine (Br₂) vapors upon exposure to light.
  2AgBr(s) --[Sunlight]--> 2Ag(s) + Br₂(g)

3. Displacement Reaction

A chemical reaction where a more reactive element displaces (removes) a less reactive element from its compound in a solution.

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

Examples of Displacement Reactions:

1. Iron in Copper Sulfate Solution:

• Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

• Observation: The blue copper sulfate solution fades to pale green (iron sulfate). A reddish-brown coating of copper metal deposits on the silvery-grey iron nail.

2. Zinc in Copper Sulfate Solution:

• Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

• Observation: The blue copper sulfate solution becomes colorless (zinc sulfate), and reddish-brown copper is deposited.

3. Lead in Copper Chloride Solution:

• Pb(s) + CuCl₂(aq) → PbCl₂(aq) + Cu(s)

• Observation: The blue-green copper chloride solution becomes colorless, and reddish-brown copper is deposited.

4. Double Displacement Reaction

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

• General Form: AB + CD → AD + CB
  (Memory Tip: An analogy is two pairs of partners (cation-anion pairs) swapping to form two new pairs.)

Examples:

1. Sodium Sulfate and Barium Chloride:

• Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s)↓ + 2NaCl(aq)

• Result: A white, insoluble solid (precipitate) of Barium Sulfate forms, while Sodium Chloride remains in a colorless solution.

2. Potassium Iodide and Lead Nitrate:

• 2KI(aq) + Pb(NO₃)₂(aq) → 2KNO₃(aq) + PbI₂(s)↓

• Result: A yellow, insoluble solid (precipitate) of Lead Iodide forms.

Precipitation Reactions

A double displacement reaction that forms an insoluble solid is also called a precipitation reaction. The insoluble solid formed is the precipitate. The examples of Barium Sulfate and Lead Iodide formation are both precipitation reactions.

Redox Reactions

A Redox Reaction is one where reduction and oxidation occur simultaneously. The name combines RED-uction and OX-idation. If one substance is oxidized, another must be reduced.

Example of Oxidation: Heating Copper in Air

When shiny, reddish-brown Copper (Cu) is heated in the presence of Oxygen (O₂), it forms a layer of black Copper Oxide (CuO). Copper has gained oxygen, so Copper is oxidized to copper oxide.

Oxidizing and Reducing Agents

Agents are identified from the reactants.

• Oxidizing Agent (Oxidant): The substance that undergoes reduction.

• Reducing Agent (Reductant): The substance that undergoes oxidation.

Example: Heating Copper Oxide with Hydrogen Gas

• Reaction: CuO + H₂ → Cu + H₂O

• Analysis:

1. Copper Oxide (CuO) → Copper (Cu): This is the removal of oxygen, hence reduction. Copper Oxide, being reduced, acts as the oxidizing agent.

2. Hydrogen (H₂) → Water (H₂O): This is the addition of oxygen, hence oxidation. Hydrogen, being oxidized, acts as the reducing agent.

Types of Redox Reactions in Daily Life

Rancidity and corrosion are common examples of redox reactions, specifically oxidation.

1. Rancidity (Vikritgandhita)

Rancidity is the oxidation of fats and oils in food, resulting in an unpleasant bad smell and taste.

Methods of Prevention:

1. Add Antioxidants: Substances that inhibit or prevent oxidation (e.g., Vitamin C and Vitamin E).

2. Replace Oxygen with Unreactive Gas: Fill food packaging with inert gas (e.g., chip packets flushed with Nitrogen (N₂) or Helium (He)) to prevent oxidation.

3. Refrigeration: Storing food at low temperatures slows down the rate of oxidation.

2. Corrosion (Sansaar)

Corrosion is the gradual degradation of a metal's surface as it reacts with atmospheric gases, converting into a more stable chemical form (e.g., oxide, sulfide, or carbonate). This typically causes the metal to lose its luster.